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LIBRARY-CHEMIS TRY 


IN 


Aluminum 
Antimony 
Argon 
Arsenic 
Barium 
Beryllium 
Bismuth 
Boron 
Bromine 
Cadmium 
Calcium 
Carbon 
Cerium 
Cesium 
Chlorine — 
Chromium 
Cobalt 
Columbium 
Copper 
Dysprosium 
Erbium 
Europium 
Fluorine 
Gadolinium 
Gallium 
Germanium 
Gold 
Hafnium 
Helium 
Holmium 
Hydrogen 
Indium 
Jodine 
Iridium 
Iron 
Krypton 
Lanthanum 
Lead 
Lithium 
Lutecium 
Magnesium 
Manganese - 


UNIVERSITY OF ILLING 


—_ ee ee 


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THE UNIVERSITY 
OF ILLINOIS 
LIBRARY 

41 93.1° Ruthenium Ru 
29 63.57 Samarium Sm 
66 162.52 Scandium Se 
68 167.7 Selenium Se 
63 152.0 Silicon Si 

Dee UU) Silver Ag 
64 157.26 Sodium Na 
31 69.72 Strontium Sr 
D2) 172.00 Sulfur S 
TO LOTR Tantalum Ta 
£2 tek io.0 Tellurium Te 

2 4.00 Terbium Tb 
67 163.4 Thallium EI 

1 1.008 Thorium Th 
49 114.8 Thulium Tm 
53 - 126.932 Tin Sn 
(WG AAS yas Titanium Ti 
26 8=6. 55. 84 Tungsten W 
a6 7 782.9 Uranium U 
57 . 138.90 Vanadium V 
82 207.20 Xenon Xe 

3 6.940 Ytterbium Yb 
TET 5S0 Yttrium me 
12 24.32 Zine Zn 
25 54.93 Zirconium Zr 


540 


At. At. . 
mber weight 
80 200.61 
42 96.0 
144.27 | 
20.2 
58.69 
14.008: 
"190.8 
16.000 
106.7 
31.027 
"195.23 
39.096 
140.92 
225.95 
222. 
102. 
85. 
101. 
150. 
45S 
79. 
28. 
107. 
22. 
87. 
32. 
181. 
127. 
159. 
204. 
232. 
169. 
118. 
47. 
184. 
238. 
50. 
130. 
173. 
89. 
65. 
91. 


OF 
44 








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LIBRARY-CHEMISTa 


4.0 
372 


Nap. 


INtRENAT IONAL “ATOMIC "WEIGHTS, 1926 


Wi VA, 


Aluminum 
Antimony 
Argon 
Arsenic 
Barium 
Beryllium 
Bismuth 
Boron 
Bromine 
Cadmium 
Calcium 
Carbon 
Cerium 
Cesium 
Chlorine — 
Chromium 
Cobalt 
Columbium 
Copper 
Dysprosium 
Erbium 
Europium 
Fluorine 
Gadolinium 
Gallium 
Germanium 
Gold 
Hafnium 
Helium 
Holmium 
Hydrogen 
Indium 
Iodine 
Iridium 
Iron 
Krypton 
Lanthanum 
Lead 
Lithium 
Lutecium 
Magnesium 


Manganese - 


Symbol number weight 


Al 
Sb 


At. At. At. At? 4s 
Symbol number weight. 

to) 4:26.97 Mercury Hg 80 200.61 
DL OI Te Molybdenum Mo 42 96.0 
18 39.91 Neodymium Nd 60 144.27 | 
33 74.96 Neon Ne 5107 220: 2 
56 1387.37 Nickel Ni 28° “358; Ges 

4 9.02 Nitrogen N 7 74.08 
83 209.00 Osmium Os 76 "190.8 

5 10.82 Oxygen O 8 16.000 
35 79.916 Palladium Pd 46> 1106.7 
48 112.41 Phosphorus 's 15° 3h 027 
20; 40.07 Platinum Pt, 78° “95s Ze 

6 12.000 Potassium K 19 39.096 
58 140.25 Praseodymium Pr 59 140.92 
55 132.81 Radium Ra,» 88 225.95 
Le 35.457 Radon Rn > 86. 222: 
DAG Oe SOL Rhodium Rh 45 102.91 
27 = 58.94 Rubidium Rb) 337) “B5c44 
4b es 9301 Ruthenium Ruse 442 | Oke tie 
29 63.57 Samarium Sm 62 150.43 
66 162.52 Scandium Seu. 2h") 450 
68 167.7 Selenium Se 34. 79.2 
6a) 15220 Silicon Si 14 28.06 

9 19.00 Silver Ag 47 107.880 
64 157.26 Sodium Naeit 22.997 
lM OO tre Strontium Sr +. 38 (87.63 
o2 72.60 Sulfur S 16 32.064 
79 197.2 Tantalum Tar? 737%) T8bh 
(oe NEE Tellurium és” 52° 12iee 

2 4.00 Terbium Th? 65. 16022 
67 163.4 Thallium Ay) 81 204.39 

1 1.008 Thorium Th +~90)- 232215 
49 114.8 Thulium Tm 69 169.4 
Dow 20. Gee) eal Sn. 350 21876 
FLA LOB TI Titanium Ti Bee CAT 
26 8655.84 Tungsten W 74 184.0 
36 82.9 Uranium U 02. 238217 
57 . 138.90 Vanadium V 23 50.96 
82. © 207.20 Xenon Xe> 54° 130:2 

3 6.940 Ytterbium Wow) 1 ¥ous 
(at 5.0 Yttrium A's 39 89. 
128 t 24232 Zine Zi" ~ 30" ' 65238 
25 54.93 Zirconium Zits 3405 391% 








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LIBRARY 
OF THE 
UNIVERSITY OF ILLINOIS 


UNIVERSITY OF ILLINOK 
LIGKARY-GHEMISTRY 


80°75 70-65 


[Pete a 


i 
Fi. | | | 
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‘i Nga 


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PLATE OF SPECTRA | 


The numbers at top and bottom of plate give wave lengths 
for each spectrum, in hundredths of a micron (2), 





A 
TEXTBOOK OF CHEMISTRY 


REVISED EDITION 


BY 


WILLIAM A. NOYES 


DIRECTOR OF THE CHEMICAL LABORATORY 
OF THE UNIVERSITY OF ILLINOIS 


UNIVERSITY OF ILLINOIS 
LIBRARY-GaEHHSray 








NEW YORK 
HENRY HOLT AND COMPANY 


COPYRIGHT, 1913, 1926 
BY 
HENRY HOLT AND COMPANY 


October, 1926 


PRINTED IN THE 
UNITED STATES OF AMERICA 





PREFACE 


Tus book is the result of twenty-five years of experience in 
teaching chemistry to classes of students, some of whom were 
beginning the subject while others have had an elementary 
course in the High School. It is intended to make the pres- 
entation sufficiently complete and clear so that college stu- 
_ dents who have had an elementary course in physics may use 
the book successfully, without any previous study of chemis- 
try. At the same time it is believed that the book is of such 
a character that the college students who use it after taking 
a course in the High School will not feel that they are wasting 
their time. It is desirable that the two classes of students 
should be separated for purposes of instruction, and it is our 
experience at the University of Ilhnois that the second class 
succeed much better than the first. We feel that the teaching 
of chemistry in the High School should be encouraged, both 
for the large class of students who go no farther and also for 
those who intend to continue the subject in college. 

The author has always tried to remember that he was writ- 
ing for students who have scarcely more than begun the subject 
and not for experienced chemists. For this purpose clear, direct 
statements are most useful, to furnish the starting point from 
which a larger knowledge may be developed later. While 
positive error should always be avoided, it is often impossible 
to secure complete, rigorous accuracy without sacrificing that 
simplicity of statement which every successful teacher knows 
to be essential. 

The atomic theory has been made the fundamental basis for 
the work and has been treated throughout as essentially true 
and not merely as a convenient hypothesis likely to be some 

lll 


CODA G2 


iV _ PREFACE 


time displaced by something very different. It is believed 
that the development of our knowledge during the last few 
years fully justifies this course. The theory of ionization has 
also been freely used, as the only means we have by which 
a large class of phenomena can be clearly presented and 
understood. 

It seems desirable to give some material which it is not 
possible to emphasize or teach thoroughly in a brief course 
and some things which are rather for reference than to be 
learned. ‘To aid teachers and students in distinguishing such 
paragraphs, they are indicated by an asterisk. This device, 
which was, I think, first used by Professor Ostwald, is better 
adapted than the use of fine print to the need of teachers who 
may wish to make a different selection. Students are earnestly 
advised to read these paragraphs, even when they are not 
expected to acquire a full knowledge of them. 

I wish to express my very sincere thanks to the following 
members of the Chemical Staff of the University of Illinois, 
some of whom have met with me almost weekly for three 
years to read and criticize the successive chapters of the book. 
The criticisms have been very valuable and helpful. C. W. 
Balke, S. J. Bates, Edward Bartow, G. D. Beal, L. L. Burgess, 
R. 8. Curtiss, C. G. Derick, Lambert Thorp, B. 8. Hopkins, 
Helen Isham (Mattill), Grinnell Jones, C. G. MacArthur, Ellen 
S. McCarthy (Foley), D. F. McFarland, D. A. MacInnes, C. F. 
Nelson, .S. W: Parr, .G: McP. Smith, Ey Ko Strachan, Ww: 
Washburn, H. C. P. Weber. 

I wish also to express my gratitude to Professor Edward W. 
Morley, Professor Julius Stieglitz of the University of Chi- 
cago, Professor J. Bishop Tingle of MacMaster University, 
and Mr. C. M. Wirick of the Crane Technical High School of 
Chicago, who have read the proofs and made many useful sug- 
gestions. JI am also indebted to Professor R. B. Moore of the 
Bureau of Mines for some valuable criticisms of the paragraph 
on Radioactivity. . 


PREFACE TO THE SECOND EDITION 


THE past ten years have been years of very rapid advance in 
some of the most fundamental concepts of chemistry. Atoms are 
no longer the ultimate, indivisible and unchangeable particles which 
they once seemed to be. Physicists and chemists are beginning to 
gain an insight into the structure of the atoms somewhat as the 
chemists of the last half of the nineteenth century learned how to 
decipher the arrangements of the atoms of organic compounds. 
The intimate connection between electrons and chemical combina- 
tion seems to be fully established. The Periodic Table has acquired 
a new meaning and the development of X-ray spectra in close 
connection with the atomic numbers has set very narrow limits for 
the possible discovery of new elements. In this edition an attempt 
has been made to present some of these new facts and theories in 
such a manner as to be interesting and profitable for students 
beginning the subject. 

Atomic numbers have been included in the tables of atomic 
weights. The older, alphabetical form has been retained in the 
table on the inside of the cover but in the table on page 10 the 
order of the atomic numbers has been used and, incidentally, 
the fact that there are only five complete periods of elements is 
shown by the dotted lines separating the periods. 

The periodic tables have been carefully revised and the relation 
of the table to atomic structures has been briefly sketched. The 
discussions of electrons have been brought into closer agreement 
with the theory of G. N. Lewis. A number of typographical cor- 
rections have been made. The author will be very grateful for 
any others which teachers using the book can send him and for 
suggested improvements in the treatment of subjects found difficult 
for beginners. 


1a rt Nae 


TW by 5 bad 
Lh ees rey iy HAO A 





CONTENTS 


CHAPTER I 
INTRODUCTION 


The Nature of Scientific Knowledge, 1.— Subdivisions of Science, 3.-~ 
Physical Sciences, 4.— Matter and Energy, 5.— Conservation of Matter and 
Energy, 6.— Pure Substances and Mixtures, 7.— Preparation of Pure Sub- 
stances, 8.— Elements and Compounds, 9.— How Pure Substances are dis- 
tinguished from Mixtures. Law of Constant Proportion, 12. — Inductive 
Reasoning, 13.— Law of Combining Weights, 13.— The Atomic Theory, 14. 
— Selection of Atomic Weights, 16.— Formulas, 16.— Composition of Pure 
Substances, 17. — Study of Chemistry, 18. 


CHAPTER II 


OXYGEN 
SymMBoL, O. ATomic WEIGHT, 16 


Occurrence, 19. — Preparation, 19.— Collection and Storage of Gases, 22. — 
Properties of Oxygen, 22.— Oxygen and Acid Properties, 23.— Combustion. 
Effect of Concentration on a Chemical Reaction, 24.— Kindling Tempera- 
ture, 24.— Heat of Combustion. Calorimeter, 25.—The Nature of Chemical 
Energy, 27.— Catalysis, 28.--Chemical Affinity, 29.— Nomenclature, 29. 


CHAPTER III 


LAWS OF GASES 
Units oF LENGTH, WEIGHT, VOLUME, TEMPERATURE, TIME AND ENERGY 


Unit of Length. Meter, 31.— Unit of Weight. Gram, 31.— Unit of Vol- 
ume. Liter, 31.— Units of Time, 32.— Unit of Temperature, 32.— Units of 
Energy. Kilogram-meter. Erg, 32.— Centimeter-gram-second System. Ab- 
solute Units, 33.— Units of Mechanical Energy, 33.— Unit of Power, 33.— 
Units of Heat, 33.— Electrical Units, 33.— Chemical Energy, 34.— Effect of 
Pressure on a Gas. Law of Boyle, 34.—Corrections for Readings of the 
Barometer, 36.—Effect of Temperature on a Gas. Law of Charles, 38.— 
Absolute Temperatures, 39.—Significance of the Absolute Zero, 40. — Deter- 
mination of the Weight of a Liter of a Gas, 40.— Graphical Representation 
of the Gas Laws, 42. — Exercises, 43. 


Vii 


vill CONTENTS 


CHAPTER IV 
HYDROGEN 


Sympou, H. Atomic WEIGHT, 1.0078 


Occurrence, 45.—Radicals, 46.— Salts, 47.—Preparation of Hydrogen. 
1, Electrolysis of Dilute Sulfuric Acid, 47.—Electrolytes. Ions. Theory of 
Electrolysis, 48.—2, Preparation of Hydrogen from Iron and Steam, 48. 
Reversible Reactions, 50.—3, Decomposition of Water by Metals at Ordi- 
nary Temperatures, 50.— Contrast between the Action of Iron and of Sodium 
on Water, 51.—4, Hydrogen from ‘“‘ Hydrone,’’ 52.—5, Preparation of 
Hydrogen by the Action of Metals on Acids, 52.— Apparatus for the Prepa- 
ration of Hydrogen, 53.—Purification of Hydrogen, 54.— Properties of 
Hydrogen, 55.— Diffusion of Gases, 56.— Kinetic Theory of Gases, 58.— 
Chemical Properties of Hydrogen, 59.— Dissociation, 59. —The Oxyhydrogen 
Blowpipe, 61.—Explosions. Catalysis, 62.— Oxidation. Reduction, 63.— 
Valence, 63.— Heat of Combustion of Hydrogen, 65. 


CHAPTER V 


WATER, HYDROGEN PEROXIDE 


Analysis, Synthesis, 66.— Qualitative Analysis and Synthesis of Water, 66. 
— Quantitative Synthesis of Water by Volume, 66.— Composition of Water 
by Weight, 68.—The Unit for Atomic Weights, 68.— Determination of the 
Composition of Water by the Use of Copper Oxide, 69.— Determination of 
the Composition of Water by Weighing Oxygen and Hydrogen, 71.— Proper- 
ties of Water, 72. —Heat of Fusion and Vaporization, 74. — Vapor Pressure of 
Water, 74.— Equilibrium, 76.— Effect of Water Vapor on the Pressure of a 
Gas, 76.— Phases. Degrees of Freedom, 77.— Water as a Solvent. Solutes, 
79.—Chemical Activity in Solutions. Metathesis, 81.— Hydrates, Deliques- 
cence, Efflorescence, 81.— Natural Waters, 82.— Purification of Water, 83.— 
Hydrogen Peroxide, 83.— Properties and Uses of Hydrogen Peroxide, 86. — 
Tests for Hydrogen Peroxide, 85.—Structure of Hydrogen Peroxide, 86. 
— Law of Multiple Proportion, 87. 


CHAPTRingY f 
AVOGADRO’S LAW. SELECTION OF ATOMIC WEIGHTS. OZONE 


Gay Lussac’s Law of Combining Volumes, 89.— Avogadro’s Law, 91.— 
Selection of an Atomic Weight, 92. — Molecules of the Elements, 93. — Gram 
Molecular Volume, 94. — Number of Molecules in one Cubic Centimeter of a 
Gas, 95. — Allotropic Forms. Ozone, 97. — Exercises, 99. 


CHAPTER VII 
CHLORINE 


SyMBoL, Cu. ATOMIC WEIGHT, 35.46. FoRMULA, Cle 


- Occurrences of Chlorine, 100.— Preparation of Chlorine. 2, By Electrol- 
ysis of Sodium Chloride, 100. — 2, Preparation by Oxidation of Hydrochloric 


CONTENTS 1X 


Acid, 100. — 3, Preparation of Chlorine by the Deacon Process, 102. — 4, The 
Weldon Process for Chlorine, 103. — Properties of Chlorine, 104. — Chlorine 
and Water. Bleaching, 106.— Chlorine Hydrate. Phases, 107.— The Heat 
of Combination of Chlorine and of Oxygen with Other Elements, 108. — Equi- 
librium in Chemical Reactions, 108.— Principle of van’t Hoff-Le Chatelier, 
111. — Effect of Water on Chlorides. JIonization, 112. — Effect of Water on 
Chlorides. Hydrolysis, 115. — Exercises, 116. 


CHAPTER VIII 


HYDROCHLORIC ACID. OXIDES AND OXYACIDS OF CHLORINE 


Hydrochloric Acid, 118. — Properties of Hydrochloric Acid, 119. — 1, Reac- 
tion with Metals, 120.— 2, Reaction with Hydroxides of Metals, 121.— 3, Re- 
action with Oxides of Metals, 122.— 4, Reaction with Oxidizing Agents, 122. 
— Indicators, 122. — Oxides and Oxygen Acids of Chlorine. Nomenclature, 
123.— Hypochlorous Acid. Hypochlorites, 124.— Hypochlorous Anhydride, 
or Chlorine Monoxide, 126.—Chlorous Acid and Chlorites, 127. — Chloric 
Acid and Chlorates, 127. — Chlorine Peroxide, 127.— Perchlorates and Per- 
chloric Acid, 128.— Structure of the Oxyacids of Chlorine, 130. — The 
Atomic Weight of Chlorine, 130. 


CHAPTER IX 
CLASSIFICATION OF THE ELEMENTS. THE PERIODIC SYSTEM, 132 


CHAPTER X 
THE HALOGEN FAMILY 


General Properties of the Halogens, 139.— Compounds of the Halogens 
with Hydrogen and Oxygen, 139.— Bromine, Br, 79.92. Occurrence. Prepa- 
ration, 140.— Properties, 141.— Hydrobromic Acid, 142.—Sodium Hypo- 
bromite, 143. Iodine, I, 126.93. Occurrence, Preparation, 144. — Properties 
of Iodine, 144. — Hydriodie Acid, 145. — Direct Combination of Hydrogen and 
Iodine. Reversible Reactions. Equilibrium, 146.—- Speed of Chemical Reac- 
tion, 148. — Concentration and Speed of Reaction, 149.— Calculation of the 
Relative Speed of Two Reactions from the Composition of an Equilibrium 
Mixture, 151.— Effect of Removing One of the Reacting Substances. Dis- 
placement of the Equilibrium Point, 152. — Heat of Formation of Hydriodic 
Acid, 152. — Flourine, F, 19.0. Occurrence, 153.— Preparation, 153. — Prop- 
erties, 154. — Etching Glass. Hydrofluoric Acid, 154. — Metallic Elements of 
Group VII, 156. — Exercises, 156. 


CHAPTER XI 


SULFUR, SELENIUM AND TELLURIUM 


Sulfur, S, 32.06. Occurrence, 160. — Allotropic Forms of Sulfur, 162. — 
Liquid Forms of Sulfur, 162. — Gaseous Forms of Sulfur, 163. — Properties 


x CONTENTS 


and Uses of Sulfur, 163.— Hydrogen Sulfide, 164.— Solution of Hydrogen 
Sulfide. Henry’s Law, 165.—Sulfides. Groups of Analytical Chemistry, 166. 
— Hydrosulfuric Acid. Strength of Acids, 167.— Application of the Idea of 
Strength of Acids to Explain the Conduct of Sulfides, 168. — Hydrogen Sul- 
fide as a Reducing Agent, 171.— Sulfur Dioxide, 172. — Sulfurous Acid, 174. 
— Sulfites, 175.— Sulfur Trioxide, 175.— Sulfuric Acid, 177. —The Electron 
Theory, 181.— Sulfuric Acid as a Dehydrating Agent, 182.— Sulfates. Di- 
basic Acids, 183,— Normal, Standard and Formular Solutions, 183. — Acid- 
imetry and Alkalimetry, 185.— Pyrosulfates, 186. — Hyposulfites, 186.— 
Thiosulfates, 186.— Persulfuric Acid, 187.— Permonosulfuric Acid, 188. — 
Polythionic Acids, 188.— Compounds of Sulfur Containing Halogens, 188. 
— Sulfur Monochloride, 188. — Chlorosulfonic Acid, 189. — Sulfuryl Chloride, 
189. — Selenium, Se, 79.2, 189. — Hydrogen Selenide, 190.— Selenium Dioxide, 
190. — Selenic Acid, 190.— Tellurium, Te, 127.5, 190.— Atomic Weight of 
Tellurium, 190.— General Properties of the Elements of the Sixth Group, — 
191. — Crystals, 192.— 172, The Isometric or Regular System, 193.— 2, The 
Tetragonal System, 194. — 3, The Rhombic System, 194. — 4, The Hexagonal 
System, 194.— 5, The Monoclinic System, 195.—6, The Triclinic System, 196. 
— Exercises, 196. 


CHAPTER XII 


NITROGEN 


Sympou, H. Atomic WEIGHT, 14.01 


Occurrence and Natural History of Nitrogen, 198. — Preparation and Prop- 
erties of Nitrogen, 200.— Ammonia, 201.— Properties of Ammonia, 202. — 
Aqua Ammonia, 203. —Ice Machines, 204. — Derivatives of Ammonia, 205. — 
The Electron Theory, 206. — Solutions in Liquid Ammonia, 207. —The Volu- 
metric Composition of Ammonia, 208.— Nitric Acid, 210.— Hydrates of Nitric 
Acid, 211.— Chemical Properties of Nitric Acid, 212. Aqua Regia, 213. — 
Nitrosyl Chloride, 214.— Oxides of Nitrogen, 214, — Nitrous Oxide, 214. — 
Nitric Oxide, 215. — Nitrous Anhydride, 218. — Nitrous Acid, 218. — Nitrogen 
Peroxide and Nitrogen Tetroxide, 219.— Nitrogen Pentoxide, 220. — Other 
Compounds of Nitrogen, 220.— Hyponitrous Acid, 221.— Hydroxylamine, 221. 
— Hydrazine, 222. — Hydronitric Acid, or Azoimide, 223. — Iodine Trinitride, 
223. — Nitrogen Trichloride and Chloroamine, 224. — Nitrogen Iodide, 225. 
— Endothermic Compounds, 225. — Exercises, 225. 


CHAPTER XIII 
THE ATMOSPHERE. NOBLE GASES 


Determination of Oxygen, 227. — Composition of Air, 227. — Air is a Mix- 
ture, 228.— Carbon Dioxide in the Air, 229. — Ventilation, 230. — Moisture, 
231. — Liquid Air. Critical Temperature, 232. — Argon, A, 39.91, 235. — 
Atomic Weight of Argon. Specific Heat of Gases, 236. — Helium, He, 4.00, 
237. — Neon, Krypton, Xenon and Niton, 238. — Exercises, 238. 


CONTENTS rierexl 


CHAPTER XIV 


PHOSPHORUS 


Phosphorus , P, 31.03. Occurrence, 240. — Preparation of Phosphorus, 241. 
— Allotropic Forms of Phosphorus, 241. — Matches, 242. — Phosphine, 243. — 
Phosphonium Salts, 243.— Phosphorus Trichloride, and Phosphorus Penta- 
chloride, 244.— Hydrolysis of the Chlorides of Phosphorus, 245.— Phosphorus 
Oxychloride, 245.— Oxides of Phosphorus, 246. — Acids of Phosphorus, 246. — 
Basicity of the Acids of Phosphorus, 247.— Hypophosphorous Acid, 248. — 
Phosphorous Acid, 248. — Orthophosphoric Acid, 248.—TIonization of Ortho- 
phosphorie Acid, 250.— Decomposition of Primary and Secondary Salts of 
Orthophosphoric Acid, 252.— Pyrophosphorie Acid, 253. — Metaphosphoric 
Acid, 253.— Hypophosphoric Acid, 254.—Sulfides of Phosphorus, 254. — 
Exercises, 255. 


CHAPTER XV 
ARSENIC, ANTIMONY AND BISMUTH 


Arsenic, As, 74.96. Occurrence, 256.— Preparation and Properties of Arsenic, 
257. — Arsine, Marsh’s Test, 257. — Arsenic ‘‘ Trioxide,’’ 258. — Arsenious Acid, 
259. — Arsenic Pentoxide and Arsenic Acid, 259.— Arsenic Trichloride, 260. 
—Sulfides of Arsenic, 260.— Arsenic Disulfide, or Realgar, 260. — Arsenic 
Trisulfide, or Orpiment, 260. — Arsenic Pentasulfide, 260. — Sulfarsenites and 
Sulfarsenates, 261.—Colloidal Arsenic Trisulfide, 261.—Antimony, Sb, 121.77. 
Occurence and Preparation, 263. — Properties, 264. — Uses, 264. — Stibine, 
264. — Oxides of Antimony, 265.— Antimony Hydroxide. Antimonious Acid, 
265. — Tartaric Emetic, 266.— Antimonic Acids, 267.— Chlorides of Anti- 
mony, 267. — Antimony Trichloride, 267. — Antimony Tetrachloride, Hydro- 
tetrachloroantimonic Acid, 267. — Antimony Pentachloride, 267. — Metachlo- 
roantimonic Acid, 268. — Antimony Trisulfide, 268. — Antimony Pentasulfide, 
268. — Sulfantimonites and Sulfantimonates, 268.— Bismuth, Bi, 209. Occur- 
rence, Properties, Uses, 268. — Oxides of Bismuth, 269.— Bismuth Chloride, 
269. — Bismuth Nitrate, 270.— Bismuth Trisulfide, 270. — Tables of Compounds 
of the Elements of the Fifth Group, 270.— Vanadium, Columbium Tantalum, 
271. — Exercises, 272. 


CHAPTER XVI 


CARBON 


Carbon. Occurrence, 273. — Diamonds, 274. — Graphite, 276. — Amorphous 
Carbon, 277. — Lampblack, 277.— Wood Charcoal, 277. — Animal Charcoal 
and Bone Black, 278. — Coke, 278. — Gas Carbon. Carbon Electrodes, 279. — 
Coal, 280. — Chemical Properties of Carbon, 281. 


CHAPTER XVII 


HYDROCARBOUS, ILLUMINATING AND PRODUCER GAS. FLAME 


Marsh Gas or Methane, 286.— Substitution, 287.— The Davy Safety 
Lamp, 287. — Homologues of Methane, 289. -— Petroleum, 289.— Ethylene or 


Xil CONTENTS 


Ethene, 290. — Unsaturated Compounds. Ethylene Chloride and Ethylene 
Bromide, 291.— Acetylene, 292.— Benzene, 294. — Illuminating Gas, 295. — 
Oil Gas, 296. — Water Gas, 296. — Producer Gas, 297. — Blast-furnace Gas, 
298. — Luminous Flames, 299. — Bunsen Burner, 300. — Explosion Waves, 301. 
— Temperature of Flames, 302. — Blowpipe, 303. — Reversed Flames, 304. — 
Exercises, 305. 


CHAT Eihex Vili. 


OXIDES AND SULPHIDES OF CARBON. ASSIMILATION AND 
RESPIRATION. CYANIDES 


Carbon Dioxide, 306.— Isothermals of Carbon Dioxide, 307. — Density of 
Carbon Dioxide, 308. — Aqueous Solutions of Carbon Dioxide, Carbonic Acid, 
309. — Carbonates and Bicarbonates. Hard Waters, 309. — Carbon Monoxide, 
311. — The Cycle of Carbon in Nature, 312. — Respiration Calorimeter, 313. — 
Carbon Suboxide, 316.—Carbon Oxychloride, or Phosgene (Carbonyl Chlo- 
ride), 316.—Carbon Bisulfide, 317.—Sulfocarbonates, 317. —Sulfocarbonic 
Acid, 318. — Carbon Oxysulfide, 318. — Cyanides, 319. — Hydrocyanic Acid, or 
Prussic Acid, 319.— Potassium Cyanide, 319.— Complex Cyanides, 319.— 
Potassium Cyanate, 321.— Potassium Thiocyanate, 321.— Cyanogen, 322.— 
Exercises, 322. 


CHAPTER XIX 


. ALCOHOLS, ALDEHYDES, KETONES, ACIDS, FATS, CARBOHY- 
DRATES 


_ Structural Formulas, 323.— 7, Valence, 323. — 2, Radicals, 323. — 3, Substi- 
tution, 324.— Alcohols, 324.— Methyl Alcohol, 324.— Ethyl Alcohol, 325. — 
Phenol or Carbolic Acid, 326.— Glycerol, 326.— Aldehydes and Ketones, 327. 
— Formaldehyde, 327.— Benzaldehyde, 328.— Acetone, 328:— Acids, 328. 
Formic Acid, 329.— Acetic Acid, 329.—Oxalic Acid, 329.—Lactie Acid, 330. 
—Tartaric Acid, 330.— Citric Acid, 330.— Ammonium Ferric Citrate, 331. — 
Benzoie Acid, 331.— Palmitic, Stearic and Oleic Acids, Fats, 331.— Soaps, 
332. — Carbohydrates, 332. — Cane Sugar, or Saccharose, 333. — Maltose, 334. — 
Lactose, or Milk Sugar, 334. — Glucose, 334. — Fructose, 335. — Starch, 335. — 
Dextrin, 336.— Pectose, Pectin, 337.— Cellulose, 337.— Paper, 337.— Gun 
Cotton, Celluloid, Lacquers, 338. 





CHAPTER XX 


AMINES, DYES, ALKALOIDS, PROTEINS, ENZYMES, FOODS AND 
NUTRITION 


Methyl Amine, 339. — Aniline, 340. — Dyes, 340.— Alizarin, 341. — Indigo, 
341. — Mordants, 342. — Alkaloids, 342. — Nicotine, 342. — Coniine, 342. — 
Atropine, 343. — Cocaine, 343. — Morphine, 343. — Quinine, 343. — Strychnine, 
343. — Ptomaines, 343. — Proteins, 343. — Enzymes, 344. — Toxins, Antitoxins, 
344. — Urea, 345. — Nutrition, 345. 


CONTENTS Xl 


CHAPTER XXI 


SILICON, BORON, GERMANIUM, TIN, LEAD, TITANIUM, 
ZIRCONIUM, CERIUM, THORIUM 


S1LIcon, Si, 28.06 


Occurrence, 348. — Preparation, 349. — Hydrogen Silicide, 349. — Silicon 
Carbide, Carborundum, 349. — Silicon Fluoride, 350. — Fluosilicice Acid, 350. 
Silicon Tetrachloride, 351. — Silicon Hexaiodide, 351. — Silicon Dioxide, or 
Silica, 351. — Artificial Silicates, 352. — Silicic Acids, 353.— Natural Silicates, 
355. — Calculation of the Formula of a Mineral, 356. — Dialysis, Semiper- 
meable Membranes, 357. — Osmotic Pressure, 358. — Germanium, 361. 
— Tin and Lead, 361. — Titanium, 362. — Zirconium, 363. Cerium, 363. 
— Hafnium, 364. — Thorium, 364. — Welsbach Mantles, 364. — Boron, 
365. — Preparation, Properties, 365. — Boron, Trioxide, Borax Beads, 365. 
— Boric Acid, 366. — Other Acids of Boron, 366. — Borax, 367. — Sodium 
Perborate, 367. — Other Compounds of Boron, 367. — Exercises, 368. 


CHAPTER XXII 


METALLIC ELEMENTS. DIFFERENCES BETWEEN METALS AND 
NON-METALS. PREPARATION OF COMPOUNDS 


Metals and Non-metals, 369.— Classification of the Metals, 370.— Melting 
Points of the Elements, 372.— Preparation of Chemical Compounds, 372.— 
Effect of Volatility, 374. Effect of Insolubility, 376.— Effect of a Common 
Ton. Solubility Product, 377.— Formation of Complex Ions, 378.— Degree 
of Ionization, 379.—Effect of Degree of Ionization, Neutralization, 384.— 
Hydrolysis, 385.—TIllustration of the Strength of Acids, 386.— Use of Indi- 
cators, 387. Systematic Study of the Metals, 390.— Metallurgy, 390.— Ox- 
ides, 392.— Hydroxides, 392.— Solubility of Salts, 393. 


CHAPTER XXIII - 


ALKALI METALS: LITHIUM, SODIUM 


General Properties of the Alkali Metals, 395.—Lithium, 395.— Lithium 
Urate, 396.— Atomic Weight of Lithium, Law of Dulong and Petit, 396.— 
The Quantum Theory, 398.—Sodium, 398.— Metallurgy, Properties, 399.— 
The Alkali Industry, 400.—Sodium Hydroxide, 401.—Sodium Oxide, 404. 
— Sodium Peroxide, 404. —Sodium Chloride, 404.— Sodium Sulfate, Glauber’s 
Salt, 406.— Acid Sodium Sulfate or Sodium Bisulfate, 408.—Sodium Sulfite, 
408. — Acid Sodium Sulfite, or Sodium Bisulfite, 408.— Sodium. Hyposulfite, 
408.— Sodium Thiosulfate, 408.—Sodium Tetrathionate, 409.— Sodium Sul- 
tide, 409.— Sodium Hydrosulfide, 409.—Sodium Nitrate, 410.—Sodium Ni- 
trite, 410.—Sodamide, 410.— Sodium Trinitride, 410.— Disodium Phosphate, 
410.—Sodium Carbonate or Sal Soda (Washing Soda). The-Leblane Soda 


X1V CONTENTS 


Process, 411.— Sodium Bicarbonate or Baking Soda, The Ammonia Soda Pro. 
cess, 412. — Sodium Silicate, or Soluble Glass, 413.—Sodium Tetraborate, or 
Borax, 413. 


CHAPTER XXIV 


ALKALI METALS: POTASSIUM, AMMONIUM, RUBIDIUM, CAESIUM. 
THE SPECTROSCOPE 


Potassium. Occurrence, 414. — Metallic Potassium, 415.— Potassium Oxide, 
415.— Potassium Hydroxide, 415.— Potassium Chloride, 416.— Potassium 
Chlorate, 416.— Potassium Perchlorate, 416. — Potassium Iodide, 417.— Potas- 
sium Polyiodides, 417.— Potassium Sulfates, 417.— Acid Potassium Sulfate, 
or Potassium Bisulfate, 417.— Potassium Nitrate, or Saltpeter, 417.—-Gun- 
powder, 418.— Potassium Nitrite, 419.— Potassium Carbonate, 419. — Potas- 
sium Bicarbonate, or Saleratus, 420. — Potassium Cyanide, 420.— Ammonium, 
420.— Ammonium Hydroxide, 420.— Ammonium Chloride, 421.— Ammonium 
Sulfide, 421.— Ammonium Hydrosulfide, 421.—Ammonium Sulfate, 422.— 
Ammonium Nitrate, 422. — Ammonium Nitrite, 423. — Ammonium Sodium 
Hydrogen Phosphate, 423.— Ammonium Carbonate, 423.— Ammonium Bi- 
carbonate, 423.— Ammonium Chloroplatinate, 423.— Rubidium and Caesium, 
424.—Spectrum Analysis, 424. 


CHAPTER XXV 


THE ALTERNATE METALS OF GROUP I. COPPER, SILVER, GOLD. 
PHOTOGRAPHY 


Copper. Occurrence, 428.— Metallurgy, 428. — Electrolytic Refining of Cop- 
per, 429.— Properties of Copper, 430.— Alloys of Copper, 431.— Copper Hy- 
droxide, 431.—Cupric Oxide, 432.—Cuprous Oxide, 432.— Cupric Chloride, 
432.— Cuprous Chloride, 432.—Cuprous Iodide, 433.—Cupric Sulfide, 483. — 
Copper Sulfate, or Blue Vitriol, 433.— Vitriols, 434. —Cupric Nitrate, 434.— 
Ammoniocupric Sulfate, 434.— Cuprous Cyanide, 434.— Precipitation of Cop- 
per by Iron, Electromotive Series, 435.— Faraday’s Law, 438.— Silver, 439. — 
Metallurgy, 439.— Pattison’s Process, 439.—Cupellation, Assaying, 440.— 
Parke’s Process, 440.— Amalgamation Process, 441.— Other Processes for the 
Recovery of Silver, 441.— Properties of Silver, Alloys, 442.—Silver Plating, 
442.— Silver Oxide, 442.— Silver Peroxide, 443. — Silver Nitrate, 444. — Silver 
Nitrite, 444.— Silver Sulfate, 444.—Silver Chloride, Silver Bromide, Silver 
Iodide, 444.— Photography, 444.—Gold, 445.— Metallurgy, 446. — Cyanide 
Process, 446.— Properties of Gold, 448.— Alloys of Gold, 448.— Oxides of 
Gold, 448. — Gold Hydroxide, 448.— Chlorides of Gold, 450. — Exercises, 450. 


CHAPTER XXVI 
GROUP II. ALKALI-EARTH METALS: BERYLLIUM, CALCIUM, 
STRONTIUM, BARIUM, RADIUM 


Beryllium, 451.—Calcium. Occurrence, 452.— Preparation, Properties, 452. 
-Calcium Hydride, 452.—Calcium Oxide, 452.— Dissociation of Calcium 


CONTENTS XV 


Carbonate and the Phase Rule, 453.—Mortar, 454.—Cement, 454.—Cal. 
cium Chloride, 455.—Chloride of Lime, 455.—Calcium Chlorate, 456.— Cal- 
cium Flouride, 456.—Calcium Sulfide, 456.— Acid Calcium Sulfite, 457.— 
Calcium Sulfate, Plaster of Paris, 457.— Plaster of Paris and the Phase Rule, 
458.—Calcium Nitrate, 460.—Calcium Phosphates, 460.—Solubility of Cal- 
cium Phosphates, 461.—Calcium Carbide, 462.—Calcium Cyanamide, 462. — 
Calcium Carbonate, 463.—Hard Waters, 463.— Determination of Free and 
Combined Carbonic Acid in Natural Waters, 464.—Calcium Acetate, 465.— 
Calcium Oxalate, 465.—Calcium Silicate, 466.— Glass, 466.— Strontium. Oc- 
currence, 467.—Strontium Hydroxide, 468.—Strontium Nitrate, 468.— Ba- 
rium. Occurrence, 468.—Barium Oxide, 468.—Barium Peroxide, 469.- 
Barium Hydroxide, 470.—Barium Chloride, 470.— Barium Nitrate, 470.—~ 
Barium Sulfide, 470.—Barium Sulfate, 470.—Flame Colors for Calcium, 
Strontium and Barium, 471.— Radium, 471.— Disintegration of Atoms, 472. 
—Nature of the Radiations from Radioactive Substances, 473.— The Life of 
an Element, 474. — Shift in the Periodic System of the Disintegration Prod- 
uct of an Element, 475.— Other Radioactive Elements, Isotopes, 475. 
— Chemical Action of the Rays, 476. — Radiochemistry in Relation to 
Geology and Medicine, 477. 


CHAPTER XXVII 


ALTERNATE METALS OF GROUP II. MAGNESIUM, ZINC, 
CADMIUM AND MERCURY 


Magnesium, 478.— Preparation, Properties, 478.— Magnesium Oxide, 479. 
— Magnesium Hydroxide, 479. — Magnesium Chloride, 480.— Magnesium Am- 
monium Chloride, 480.— Magnesium Sulfate, 480.— Magnesium Sulfide, 480. 
— Magnesium Ammonium Phosphate, 480.— Zine. Occurrence, 481.— Metal- 
lurgy, 481.—Uses. Galvanized Iron, 481.—Sherardized Iron, 482.— Zine 
Oxide, 482.— Zinc Hydroxide, 483.— Zine Chloride, 483.— Zine Sulfate, or 
White Vitriol, 483.— Zinc Sulfide, 483. — Cadmium, 488.— Cadmium Hydrox- 
ide, 484.— Cadmium Sulfate, 484.— Cadmium Sulfide, 484.— Mercury, Hg, 200.6. 
Occurrence. Metallurgy, 484.— Properties and Uses, 485.— Amalgams, 486. — 
Compounds of Mercury, 488.— Mercurous Oxide, 488. — Mercuric Oxide, 488. 
Mercurie Sulfide, 489.—Mercurous Chloride, or Calamel, 489.— Mercuric 
Chloride, or Corrosive Sublimates, 489.— Mercuric Iodide, 490.— Mercurous 
Nitrate, 490.— Mercuric Nitrate, 490.— Mercuric Cyanide, 490. — Mercuric 
Fulminate, 491.—Ionization of Compounds of Cadmium and Mercury, 491.— 
Solubility of the Sulfides of Group II, 491.— Conduct of Solution of Mag- 
nesium, Zinc and Cadmium Salts towards Ammonium Hydroxide, 491. — 
Ammono-mercuric Compounds, 492.— Nessler’s Reagent, 492. — Exercises, 493. 


CHAPTER XXVIII 


METALS OF GROUP III. ALUMINIUM FAMILY. RARE EARTH 
METALS 


Aluminium, 494.— Metallurgy, 495.— Properties of Aluminium, 497.— 
Alloys, 497. — Goldschmidt’s Thermite Process, 497. — Aluminium Chlo- 
ride, 498. 


Xv1 CONTENTS 


— Aluminium Fluoride, 499.— Aluminium Hydroxide, 499.— Aluminium 
Oxide, 499.— Aluminium Sulfate, 500.— Alums, 500.— Brick, Earthenware, 
Porcelain, 501.— Ultramarine, 502.— The Rare Earths, 502.— Scandium, 503. 
Yttrium, 503. — Lanthanum, 503.— Ytterbium, 504.— Praseodymium and Neo- 
dymium, 504.—Samarium, 505.—Europium, Gadolinium, Terbium, 505.— 
Holmium, 505.—Dysprosium, 505.—Erbium, 505.—Thulium, 506.— Lute- 
cium, 506.— Gallium, 506.— Indium, 506. — Thallium, 507.— Exercises, 507. 


CHAPTER XXIX 


TIN. AND LEAD 


Tin. Occurrence, Metallurgy, 508.— Uses of Tin, Alloys, Tin Plate, 509.— 
Compounds of Tin, 510.—Stannous Oxide, 510.—Stannous Chloride, 510.— 
Stannous Sulfide, 510.—Stannic Oxide, 510.—Stannic Acids, 511.—Stannic 
Acid, 511. — Metastannic Acid, 512.— Parastannice Acid, 512.— Stannic Chlo- 
ride, 512.—Stannic Sulfide, 512.— Fireproofing of Cotton Goods, 513.— Lead. 
Occurrence, Metallurgy, 513.— Properties and Uses of Lead, Alloys, 514.— 
Oxides of Lead, 515.— Lead Monoxide, or Litharge, 515.— Storage Batteries, 
516.— Lead Sulfide, 518.—Lead Chloride, 518.— Lead Tetrachloride, 518.— 
Lead Sulfate, 519.— Lead Nitrate, 519.— Lead Acetate, or Sugar of Lead, 519. 
— Basic Lead Acetates, 519.— Lead Carbonate, 519.— Basic Lead Carbonate, 
or White Lead, 520. 


CHAPTER XXX 


VANADIUM AND CHROMIUM GROUPS 


Group V. Vanadium, 522.—Columbium (or Niobium), 523.— Tantalum, 
523.— Group VI. Chromium, 524. — Metallurgy, Uses, 524. — Chromous 
Chloride, 525.—Chromic Oxide, 525.— Chromic Hydroxide, 525.— Chromic 
Chloride, 525.— Hydrates of Chromic Chloride, 525.— Potassium Chromium 
Sulfate, or ‘‘Chrome Alum,” 527.— Potassium Chromate, 527. — Potassium 
Dichromate, or Pyrochromate, 527.— Lead Chromate, or Chrome Yellow, 527. 
— Barium Chromate, 528.— Chromium Trioxide, or Chromic Anhydride, 528. 
Chromy] Chloride, 528. — Molybdenum, 528. — Molybdium Trioxide, or Molyb- 
dic Anhydride, 528. —Compounds of Molybdenum, 529. — Molybdic Anhydride, 
529. — Tungsten, 530. — Compounds of Tungsten, 531. — Phosphotungstic Acid, 
531.— Uranium, 531. 


CHAPTER XXXI 


MANGANESE 


Group VII. Manganese, 533. — Occurrence, Properties, 533. — Compounds 
of Manganese, 534. — Manganous Manganic Oxide, 534.— Manganous Hydrox- 
ide, 535.— Manganous Chloride, 535.— Manganous Sulfate, 535. — Manganous 
Sulfide, 535.— Manganese Dioxide, or Black Oxide of Manganese, 535.— 
Manganates, 536.— Permanganates, 537. — Potassium Permanganate, 537.— 
Manganese Heptoxide, or Permanganic Anhydride, 538. 


CONTENTS XVll 


CHAPTER XXXII 


IRON, COBALT, NICKEL 


Group VIII. Iron, 539.— Occurrence of Iron, 540. — Metallurgy and Iron, 
540.—Pig Iron, Cast Iron, 543.— Wrought Iron, 544.—Cementation Steel, 
Cast Steel, 545.—Bessemer Steel, 547.— Open Hearth, or Siemens-Martin 
Process, 548.— Alloy Steels, 552.— Compounds of Iron, 552.— Potassium Fer- 
rate, 553. — Ferrous Chloride, 553.— Ferrous Hydroxide, 553. — Ferrous Oxide, 
553. — Ferrous Sulfate, Green Vitriol, or Copperas, 554.— Ferrous Carbonate, 
554. — Ferrous Chloride and Nitric Oxide, 554.— Ferric Chloride, 554. — Ferric 
Hydroxide, 555.— Dialyzed Iron, 555.—Ferric Oxide, 555.—Ferric Sulfate, 
556.— Magnetic Oxide of Iron, 556.—Ferrous Sulfide, 556.—Ferric Sulfide, 
556.—TIron Disulfide, or Iron Pyrites, 556.—Ferric Thiocyanate, 556.—Co- 
balt, 557. Compounds of Cobalt. Oxides, 557. — Cobaltous Hydroxide, 557. — 
Cobaltous Chloride, 557.— Cobalt Sulfide, 558.— Cobalt Nitrate, 558. — Cobalt 
Glass, 558.— Potassium Cobaltocyanide, 558.— Potassium Cobalticyanide, 
558. — Potassium Cobaltinitrite, 558.— Cobalt Ammines, 559. — Nickel, 559. — 
Compounds of Nickel, 560.— Nickel Dimethylglyoximine, 560.— Nickel Car- 
bonyl, 561. 


CHAPTER XXXIII 


THE PLATINUM METALS 
Ruthenium, 563.—Rhodium, 563.— Palladium, 563.—Osmium, 564.— 


Iridium, 565.—Platinum, 565.—Platinous Chloride, 565.—Chloroplatinic 
Acid, 565. — Platinic Chloride, 566.— Platinum Disulfide, 566. 


INDEX ‘ F F ; A Ns : ‘ 5 A ; : “ears 





A TEXTBOOK OF CHEMISTRY 


CHAPTER I 
INTRODUCTION 


The Nature of Scientific Knowledge. The phenomena pre- 
sented to our senses are so complex and varied that a complete 
description of each of them is impossible. It is the purpose 
of science to classify these phenomena and to discuss relation- 
ships between them which are frequently repeated and separate 
them from those relationships which are not repeated and which 
are to be considered as more or less accidental. When a 
given relationship between two or more phenomena is found 
always to exist, we conclude that the relationship js necessary 
or inherent in the nature of things, and it is called a law. The 
original thought conveyed by the name was, doubtless, that 
the material universe acts as it does because it was commanded 
to do so by some higher power, but the word has come to 
signify simply a statement of some constantly recurring rela- 
tionship between phenomena. 

To illustrate: we find that if we let go of an object held in 
the hand it will fall. The first time this is observed it is 
simply a fact of experience or observation. But a further 
examination of the relations involved leads to the general 
statement that any body which is not supported will fall. 
We may, perhaps, call this a law, but it is still a comparatively 
imperfect statement of the existing relations. A further study 
teaches us that all bodies fall, in a vacuum, at the same rate, 
irrespective of their size or weight and that the velocity of 
a falling body varies as the time during which it has fallen. 


These may be called the empirical laws of gravitation, that is, 
1 


2 A TEXTBOOK OF CHEMISTRY 


the laws derived from experiment and observation. For the 
class of phenomena to which they apply such laws have, prac- 
tically, as great a certainty as any individual fact which we ob- 
serve. A further study of the matter and especially of the 
motions of the earth and planets and stars has led to the more 
complete generalization that all bodies attract each other directly 
as the product of their masses and inversely as the squares of 
their distances. This is Newton’s law of universal gravitation. 
When once stated, it is seen that the empirical laws stated above 
follow from it. 

A natural law, when discovered, usually requires further study 
and development in two directions. It calls, in the first place, 
for a careful examination of the logical consequences which follow 
from the law, and it furnishes the means of predicting in count- 
less cases just what will occur in given conditions where the law 
applies. Thus the law of gravitation is assumed, instinctively, 
every time that we move. It enables us to predict, accurately, 
the motion of a pendulum or projectile and to calculate the exact 
relative positions of the sun, moon and planets for a hundred 
or a thousand years to come. The law lies at the very basis 
of all of those calculations of the engineer by which he deter- 
mines the necessary strength for the parts of a bridge or a truss. 
Many other illustrations might be given of its practical impor- 
tance and usefulness. 

In the second place, a natural law suggests the need of some 
additional explanation. An inquiry in this direction may lead 
to some more fundamental and general law, as when the laws of 
falling bodies were found to be only special cases of the universal 
law of gravitation, or it may lead to the confines of our present 
knowledge at a point where no further progress can be made 
without the use of speculation or hypothesis. Following our 
illustration, it appears to most minds almost or quite incon- 
ceivable that one body should act upon another at a distance 
without some medium, and there have been many speculations 
with regard to some medium which may be the cause of gravita- 
tion. Several hypotheses with regard to the mechanism of the 


INTRODUCTION 3 


action of such a medium have been proposed, but these specula- 
tions have not met with any notable success. 

With regard to such speculation there are, at present, among 
scientific men, and notably among chemists, two somewhat dis- 
tinct schools or classes. One of these takes the ground that the 
number of possible explanations of those parts of the universe 
which lie beyond our knowledge is so great that it is hopeless 
to attempt to find the true explanation. While admitting the 
value of hypotheses in stimulating and directing research, this 
school claims that such hypotheses can never give us any real 
knowledge of matters which are beyond the cognizance of our 
senses, and that all genuine scientific advance consists in giving 
a fuller description of things about which we can gain direct, 
positive knowledge. The other school points out that, while 
there are many things in the universe which must always remain 
beyond the possibility of direct knowledge, we can accumulate 
so much evidence with regard to these that it may be possible, 
ultimately, to give to our theories with regard to them a high 
degree of probability. The danger of the first attitude of mind 
is that the investigator will be content with a full description of 
phenomena and will fail to discover relations which can be under- 
stood only by a knowledge of matters about which we can secure 
only indirect evidence. The danger of the other point of view 
is that it may lead one to overestimate the amount of knowledge 
which has been acquired about unseen things and to spend time 
in useless speculations which would be better spent in acquiring 
new facts. 

Whichever view is accepted, the science of our time includes 
a knowledge of a very great number of facts, of the natural laws 
which express the relations between these facts and of the 
theories which are our best present explanation of the laws. 

Subdivisions of Science. There is, properly speaking, only 
one science, which includes all classified, systematic knowledge. 
The amount of such knowledge has become so great, however, 
that it is customary to subdivide it into a number of parts, each 
of which is called a science. It should always be remembered 


4 A TEXTBOOK OF CHEMISTRY 


that the boundaries between various divisions are more or less 
arbitrary and that very many facts belong about equally to two 
or more of the sciences. It is also very important to understand 
that no one can make much progress in any one science without 
considerable knowledge of several others. 

The more important subdivisions of science are the following: 
abstract sciences, which deal, primarily, with forms of abstract 
reasoning — mathematics and logic; physical sciences, dealing 
with the phenomena of matter and energy apart from life — 
physics, chemistry, astronomy, mineralogy, geology; biological 
sciences, dealing with the phenomena of living bodies — bacteri- 
ology, botany, zoology, paleontology; psychological sciences, 
dealing with the phenomena of mind and of society — psychology, 
language, history, social science, political economy, ethics. 
| Mathematics 
| Logic 

Physies 

Chemistry 
Physical sciences Astronomy 
Mineralogy 
Geology 
Bacteriology 
Botany 
Zoology 
Paleontology 
Psychology 
Language 
History 
Social science 
Political economy 
Ethics 

Physical Sciences. The two fundamental physical sciences 
are physics and chemistry. The other three mentioned, astron- 
omy, mineralogy and geology, are concerned with the applica- 
tion of the laws of physics and chemistry in studying particular 
bodies or substances and so are more special in their nature. 


Abstract sciences 


Biological sciences 


Psychological sciences 


INTRODUCTION 5 


Roughly speaking, physics treats of energy and chemistry of 
matter. Thus chemistry tells us of the properties and composi- 
tion of substances, as of water, of iron or of sulfur, of the action 
of substances on each other and of the changes in composition 
which they undergo in a great variety of circumstances. Phys- 
ics, on the other hand, deals with the varieties of energy, as 
mechanical energy, sound, heat, light, electricity, and with the 
transformations of each of these into other forms. 

As we can have no knowledge of matter except through the 
energy which it possesses and the effect of that energy on our 
senses In one way or another, and since the changes in energy 
which result when substances act on each other are often of great 
importance, the chemist can make little progress in his study 
without a considerable knowledge of physics. And as we have 
no knowledge of energy apart from matter, the physicist finds 
-some knowledge of chemistry desirable. This interrelation be- 
tween the sciences is so close, also, that there is a large domain 
which is common to both and of which it is scarcely worth while 
to ask whether it belongs to chemistry or to physics. 

Matter and Energy. The two most fundamental concepts of 
physical science are matter and energy. Matter may be defined 
as anything which has mass, or, in its relation to the earth or to 
other bodies, weight.!. Putting the same thought in quite different 
words, matter is anything which requires energy to set it in mo- 

1 This definition of matter, which is based, of course, on Newton’s 
conception of inertia — that no body can move, if at rest, or change 
the direction or velocity of its motion, if moving, unless it is acted 
on by some external force —is not entirely satisfactory. It has 
been shown that the mass of a body is changed by a change in its 
velocity, though the change is inappreciable until the velocity 
approaches that of light. (D. F. Comstock, J. Amer. Chem. Soc., 30, 
683). In such a case it is usually better to retain the older, simple 
definition, frankly recognizing that it is imperfect. It is character- 
istic of a scholastic rather than a scientific attitude of mind to be 
much troubled because a definition is imperfect or incomplete. The 
scientific worker sees that our knowledge is incomplete in every 
direction and is constantly developing. The definitions are simply 
a means of conveying this incomplete knowledge to others and of at- 
tempting to discover the most fundamental conceptions of selence. 


We eae best in both directions by keeping these definitions 
simple 


6 A TEXTBOOK OF CHEMISTRY 


tion or to change its rate of motion. Energy is usually defined, 
on the basis of etymology, as anything which can do work. A 
more satisfactory definition is that energy is anything which may 
set matter in motion or change its rate of motion. The most 
important forms of energy are mechanical energy, sound, heat, 
light, electrical energy and chemical energy. 

Conservation of Matter and Energy. Asuperficialobservatign 
of many phenomena in nature appears to show that under some 
conditions, especially in burning, matter is destroyed, and that 
under other conditions, as in the rusting of iron, matter in- 
creases in weight. It is comparatively easy to show, however, 
that a part of the air, which has weight, takes part in these pro- 
cesses and that while a candle, for instance, seems to bedestroyed 
in burning, water and carbon dioxide are formed by its combus- 
tion; andif these are absorbed by soda lime and weighed, the sum 
of their weights is very considerably greater than the weight of 
the candle which has been burned. Still more careful experi- 
ments will show that the products of combustion weigh exactly 
the same as the weight of the candle and the weight of the por- 
tion of the air (oxygen) with which it has combined. 

The question whether there is any change in the weight of 
matter during a chemical reaction is so fundamental that one 
chemist (Landolt) has considered it worth while to give ten years 
of most careful and painstaking work to its study. His con- 
clusion is that in the cases which he studied no change so great 
as the one millionth part of the weight of the substances which 
reacted with each other occurred. We say, therefore, that no 
method is known by which we can create, or destroy, matter. 
This is known as the law of conservation, or indestructibility of 
matter. 

If we place a wheel with vanes in a can of water and wind 
around its axle a cord tied to a weight in such a manner that as 
the weight falls the wheel will revolve and stir the water, we shall 
find that the temperature of the water will rise. If the experi- 
ment is carefully performed, it will be found that a weight of 
one kilogram falling 427 meters will raise the temperature of a 


INTRODUCTION i 


kilogram of water one degree. On the other hand, if the steam 
from a boiler is caused to drive the piston of a steam engine 
which is pumping water or doing other work, it is found that a 
part of the heat of the steam disappears and that the heat which 
can no longer be found in the exhaust steam or anywhere about 
the engine corresponds accurately to the amount of work which 
the engine performs, and that the ratio is exactly the same as 
that found between the falling weight and the rise in tempera- 
ture of water stirred by the paddle wheel. The energy of the 
engine may be used to drive a dynamo which will furnish an elec- 
tric current; the electric current may be used to decompose a 
chemical compound, giving substances which contain more chem- 
ical energy than the compound; and these substances, in turn, 
may be recombined, giving out heat in the process. Each form 
of energy which we know may be transformed into some other, 
and there is always an exact relation between the quantity of 
energy of one kind which disappears and the quantity of other 
kinds of energy which takes its place. This is the law of the 
conservation of energy. It might be called the law of the inde- 
structibility of energy. 

Pure Substances and Mixtures. As hasbeenstated, chemistry 
deals, primarily, with the properties and composition of sub- 
stances.! If we examine certain substances, such as pure water 
or gold, we find that they are alike throughout their whole mass, 
or that they are homogeneous. We find that every sample of 
such a substance which we examine melts or freezes at exactly 
thesame temperature ; and that if it boils without decomposition, 
it will always boil at the same temperature under atmospheric 
pressure. We find, too, that if the substance is a liquid or gas, 
the density or specific gravity is always the same under the same 
condition of temperature and pressure. In spite of the large 


1The distinction between the words body and substance should 
be carefully observed. Body always refers to some definite, con- 
crete thing, as a heavenly body, speaking of the sun or a star, 
a body of ore, ete. Substance, on the other hand, refers to some par- 
ticular kind of matter, as water or gold. <A given piece of gold 
might be called a body, but gold, in general, is a substance. 


lS 


8 A TEXTBOOK OF CHEMISTRY 


number of substances which are known to exist (more than one 
hundred thousand) it is possible to identify many of these with 
practical certainty by the examination of a comparatively small 
number of their properties. For instance no other substance has 
the same freezing point, boiling point and density as water. 
Contrasted with pure substances, as water or gold, most sub- 
stances which we meet in daily experiences are mixtures. For 
example, if we take a cereal, as wheat, and powder it, as is done 
in the milling process, a portion will pass through fine bolting 
cloth, while another portion, the bran, will not. If the portion 
which passes the bolting cloth is warmed gently, it loses weight, 
and it can readily be shown that the loss is almost wholly due 
to the escape of water, which may be condensed and identified 
by its freezing point and boiling point. From the dry flour ether 
will dissolve an oil or fat which will be left behind on evaporating 
the ether. If the portion which remains is kneaded between the 
fingers in a stream of running water, a fine white powder, consist- 
ing mainly of starch, will be washed away, while a residue, called 
gluten, which consists largely of proteins, will remain. The 
processes described show clearly that the cereal is a very complex 
mixture of many different substances, but of the substances sepa- 
rated only the water and starch can be considered as even approx- 
imately pure substances. A more careful examination of the bran 
or oil or gluten will show that each of these is still a mixture. 
Preparation of Pure Substances. A large part of the work 
which must be done in the study of chemistry consists in the 
separation and characterization of pure substances. The most 
common means used for this purpose are treatment of mixtures 
with solvents, crystallization and distillation. Thus if we have 
a mixture of sugar and sand, the sugar may be easily separated 
by dissolving it in water and pouring off or filtering the solution 
from the sand. From a brine which contains other substances 
in solution along with salt, the salt may be obtained nearly pure 
by evaporating it till the salt separates in crystals. For some 
reason particles of the same kind separate from a solution on 
evaporation or, frequently, on cooling a hot solution, in definite, 


INTRODUCTION 9 


geometrical forms called crystals, and when they separate in 
this manner they usually exclude other substances which may 
be present. By repeated distillation of a mixture of alcohol and 
water, collecting the lower boiling portions by themselves each 
time, nearly pure alcohol can be separated from water. Whena 
volatile substance contains a nonvolatile one in solution, the 
separation by distillation is much easier and 
more complete. 

Elements and Compounds. If the red 
oxide of mercury is heated in a small tube, 
metallic mercury will distill away, while a 
glowing splinter held at the mouth of the 
tube will burst into flame. The heat causes 
the decomposition of the oxide of mercury 
into mercury and a gas which supports com- 
bustion better than air, and which is called 
oxygen. An electric current passed between 
two strips of platinum immersed in a solution 
of sulfuric acid in water in the apparatus 
shown in Fig. 1 will cause the separation of 
two gases, oxygen and hydrogen. As it can 
be shown that the amount of sulfuric acid 
remains unchanged, it is evident that the 
gases are formed by the decomposition of Fig. 1 
the water; and this view can be confirmed 
by burning the mixture of oxygen and hydrogen and regenerat- 
ing the water. While oxide of mercury can be decomposed 
into mercury and oxygen, and water may be decomposed into 
oxygen and hydrogen, no one has thus far succeeded in decom- 
posing mercury or oxygen or hydrogen. Substances like these, 
which it has not been found possible to decompose, are called 
elements! Substances which can be separated into two or 





1This definition is not wholly satisfactory, since it has been 
found that radium, which has all of the other properties of an ele- 
ment, decomposes spontaneously into helium and a whole series 
of other elements. It seems best, however, to retain the simple 
definition, but also it is best to consider radium as an element. 


10 A TEXTBOOK OF CHEMISTRY 


more parts, neither of which can be converted into the other, or 
which can be prepared by the union of two or more elements, 
are called compounds. Elements sometimes exist in two or 
three different forms, but these may always be converted each 
into the other. 

Eighty-seven elements have been positively identified. There 
are reasons for believing that only five more can be discovered. 
The list is arranged in periods (p. 132). 





At. At. At. At. 
number Symbol weight number Symbol weight 
1 Hydrogen... H 1,008 © <36°Krypton™ Soi kr mis26 
2 Helium!) 2.60 elle 4.00 
aé Rubiditun sa. 85.44 

® lathiim “a. Ubi 6.940 38 Strontium .. Sr 87.63 
4 Beryllium .. Be 9.02 39_7Y Curiae ee 89. 
buBoron <s ib 10.82 40 Zirconium . . Zr 91. 

Gr Carbone earn. eke 12.000 41 Columbium . Cb 93.1 
*uNitrogen Me eae 14.008 42 Molybdenum Mo _ 96.0 
8 Oxygen ... O 16.000 43 

0 Fluorine...) 4 19.00 44 Ruthenium . Ru _ 101.7 
TOINe6on ee Ape eee 20.2 45 Rhodiim  % yen eles 

46 Palladium . . Pd _ 106.7 
Eissodium. -. 6: 2° Nat :22i997 F471 Silver’. i) a. eA eee S80 
12 Magnesium . Mg = 24.32 48 Cadmium .. Cd _ 112.41 
IS@Muminiumy, = Al 26.97 49 Indium ... In ~ 114.8 
14°Siliconl, 2 A 28.06 5027 1ns OA ae een pe 
15 Phosphorus . P 31.027. . 51 Antimony... »Sb 121.77 
16 Sulphur oes 32.064 52 Tellurium .. Te 127.5 
17 Ghiorinewesa Cl 35145790 53 *lodine aan eee 126.932 
1S-Argonwee. Ware rok 39.91 54 Xenon ... Xe _ 180.2 
19 Potassium .. K 839.096 55 Cesium “/.+/)'%CGs 7132.81 
20 Calcium ren (8) 40.07 06. Barlumagiie - oso aes 
21 Scandium .. Se 45.10 57 Lanthanum . La_ 138.90 
C2 itaniiieeee ea 47.9 58 Ceriuny (oy 29 Ce- 14025 
29) Vanadiumeen. AV" 50.96 59 Praseodymium Pr _ 140.92 
24 Chromium. . Cr 52.01 60 Neodymium . Nd _ 144.27 - 
5 25 Manganese. . Mn 54.93 61 Hinge eae 

26 ronghe = Acne Cae 55 eat 62 Samarium .. Sm _ 150.43 
Zi hvOnalt he nae eo 58.94 63 Europium) « ~ Ha 152.0 
28 Nickel™ “2°. aeN1 58.69 64 Gadolinium . Gd _ 157.26 
29 Copper’. |. i eGus 263.57 65 Ter brunisgaee De al 
DAL p A cme iy an ein’ 65.38 66 Dysprosium . Ds_ 162.52 
oa Graliume Ge oe 69.72 67 Holmium .. Ho 163.4 
32 Germanium . Ge 72.60 6S, Hirbiuing eee nn ub aden, 
SsPATSCHION. We.) on ES 74.96 69 Thulium. .. Tm 169.4 
34 Selenium .. Se 79.2 COPY tlerbiuinin eons Da. 


So wSsTOMmINGi ek ot Or 79.916, 71. Lutectum —.,. . Lu. 175.0 


INTRODUCTION Bi 


At. At. At. At. 
number Symbol weight number Symbol weight 
PP paethbeven views open bi PNW ge ae 83 Bismuth. . . Bi 209.00 
Poulan Hye and Lol. 54 9Polonium’ << «Po, 210. 
74 Tungsten .. W _ 184.0 85 
75 SOsaAcdOn st. PRIAIe eee: 
76 Osmium.) )/o..)+Os © 190.8 
Mie TIN ease LE 193.1 87 
Por igiiniuniiess 2 ott eet lon.25 Ss aan | Ay erase 225.05 
OTOL Al eee AU) 3197.2 SOsANetimiuM bhiay GAcAT 227. 
pUliMercury. .., He 200.61 CURL OOTUn eke k end eee sog. Lo 
Siernaliim: - 2. Tl 204.39 OPV mere yore keh eas 
moe leaders tee eb 220720 62; Uranium )'4 >. 1.U 238.40 


The symbols are either the first letter or the first letter together 
with some other characteristic letter of the name of the element. 
With few exceptions symbols are derived from the English names 
and the symbols readily suggest the names. The exceptions are: 


Antimony, Sb, Stibiwm Potassium, K, Kaliwm 

Gold, Au, Aurum | Silver, Ag, Argentum 
Iron, Fe, Ferrum Sodium, Na, Natriwm 
Lead, Pb, Plumbum (bay, Sn, Stannum 
Mercury, Hg, Hydrargyrum Tungsten, W, Wolfram 


For all of these except the last the symbols are derived from the 
Latin names. 

The elements vary greatly in their relative abundance. Of that 
portion of the earth which we are able to examine it is estimated 
that oxygen forms nearly one half of the total weight and sili- 
con one fourth. The percentage amounts of the twelve most 
common elements in the surface of the earth to a depth of ten 
miles, including the ocean and the atmosphere, are estimated as 
follows : 1 


PER CENT PER CENT 

Oxygen, 49.78 Potassium, 2.28 
Silicon, 26.08 - Magnesium, 2.24 
Aluminium, 7.34 Hydrogen, 0.95 
Tron, AED Titanium, 0.37 
Calcium, 3.19 Chlorine, 0.21 
Sodium, Hino Carbon, p09 

99.07 


1K, W. Clarke, Data of Geochemistry, p. 32. 


12 A TEXTBOOK OF CHEMISTRY 


Some elements which form only a very small part of the whole 
are very important, especially nitrogen, phosphorus and several 
of the metals which are not included in the above table. 

How Pure Substances are distinguished from Mixtures. 
Law of Constant Proportion. A very large part of our knowledge 
of chemistry depends on the preparation of pure substances and 
on the determination of the properties and composition of these. 
It is, therefore, important to understand how we may distinguish 
between pure substances and mixtures. The first characteristic 
of a pure substance is that it must be homogeneous so long as 
it exists in one state of aggregation, that is, so long as it is all 
solid, all liquid or all gaseous. Second, it must have a constant 
melting point and boiling point, if it melts and boils without de- 
composition, and the specific gravity or density and other physi- 
cal properties must be invariable under the same conditions.1 
Third, a pure substance must always show the same conduct 
toward any other substance which may dissolve it or act upon 
it chemically, provided that the conditions are the same. 

A very careful examination of a large number of substances 
which have the characteristics Just given in the highest degree 
has demonstrated that such substances are absolutely constant 
in composition. This is the law of constant proportion and may 
be stated thus: A pure substance always contains the same ele- 
ments in the same proportion by weight. Thus pure water always 
contains hydrogen and oxygen in the proportion of 1 to 7.94 
parts by weight. This law has been tested by a large amount 
of most careful and painstaking work, and the more careful the 
work has been the more accurately has the law been found true, 
so that we may consider it as one of the most absolutely perfect 
laws of nature. Since a very large number of substances which 
fulfill the first three requirements of a pure substance are in- 
variable in composition, this constancy of composition is con- 
sidered as a fourth characteristic of a pure substance. It is a 


_1 The density of some solids and especially of metals may vary 
money according to the treatment to which they have been sub- 
jected. 


INTRODUCTION 13 


characteristic of very great importance and one which is fre- 
quently used to determine whether a given substance is pure or 
not. 

Inductive Reasoning. It may seem at first that the use of 
constancy of composition as a means of determining whether a 
substance is pure or not is due to reasoning in a circle, or, as it is 
commonly called, is ‘ begging the question.” We say first that 
a pure substance has a constant composition and then that 
because a substance has a constant composition it is pure. The 
criticism would be justified if constancy of composition were 
the only characteristic applied to decide whether a substance is 
pure or not. But the first three characteristics mentioned above 
are the ones which will appeal to any one as being dictated by 
common sense. When we find that a very great number of 
substances having these characteristics are also constant in com- 
position, we come to the conclusion that there is some inherent, 
necessary connection between this fourth characteristic and the 
other three, and that when a given substance does not have this 
characteristic it probably lacks some of the other three as well. 
Such a conclusion is said to be reached by inductive reasoning. 
The truth of such a conclusion can never be absolutely proved 
any more than we can prove that the sun will rise to-morrow 
morning. But we may reach practical certainty by means of 
such conclusions and may properly use them as the basis for 
further reasoning. 

Law of Combining Weights. If we select a series of com- 
pounds in such a manner that each compound has an element 
contained in the preceding and another contained in the follow- 
ing compound, it will be found that whenever the same element 
recurs the proportion of the element which combines with other 
elements will always be the same or some exact multiple or 
submultiple of the first proportion. This will be more clear 
from the following series of compounds :! 


1 Here and elsewhere whole numbers are used for greater sim- 
plicity. The exact values will be found in the table of atomic 
weights, p. 10. 


14 A TEXTBOOK OF CHEMISTRY 


Water Cuprous Cupric Hydrogen Hydrochloric 
Oxide Sulfide Sulfide Acid 
Heo O: Cu Cir 5: H jis 
iP eves 8 : 63.6 63.6: 32 Oona ae 
Ferrous Ferrous Sulfur Sodium Sodium Sodium 
Chloride Oxide Dioxide Sulfide Chloride Chlorate 


Cl:Fe Fe:O O:S S:Na Na: Cl Nae Cl 
RleoGee poGeelo iis The 165.23. 23°35.) 0 wcossoDDriee 


In this series of compounds hydrogen has been chosen as the 
starting point and has been given a value of 1. If oxygen had 
been chosen and had been given a value of 100, as was at one 
time proposed, the other numbers would all be different but 
exactly the same ratios between the different numbers for the 
same element would be found throughout the series. 

It is seen that the values for hydrogen are 1 and 2, for oxygen 
8, 16, 32, and 48, for sulfur 32 and 16, for chlorine 71 and 35.5, 
the larger numbers for each element being in every case 
exact multiples of the smallest number for the element. 
This table might be extended to include all pure substances which 
have been analyzed. The law of combining weights stated 
above may be expressed more briefly as follows: A number 
may be selected for each element which represents the proportion 
of the element which enters into combination with other elements. 

The Atomic Theory. The laws of constant proportion and 
of combining weights find a very satisfactory explanation in the 
atomic theory, which was proposed by Dalton at the beginning 
of the nineteenth century. According to this theory the chemi- 
cal elements are composed of very small particles or atoms, the 
atoms of the same element being all alike in properties and in 
weight,! while the atoms of different elements are different. If 
we suppose further that compounds are always formed by the 
union of atoms of different elements, it is evident that the ratio 
between the weights of the elements in a compound must be 
the same as the ratio between the weights of the atoms composing 
the smallest particle of the compound. ‘Thus if the smallest 
particle (molecule) of water which can exist contains two atoms 


1 See, however, isotopes ( p. 475 ). 


INTRODUCTION 15 


of hydrogen united to one atom of oxygen and an atom of oxygen 
weighs 16 times as much as an atom of hydrogen, any quantity 
of water, whether large or small, must contain hydrogen and 
oxygen in the proportion of two to sixteen. If, for instance, 
1000 atoms of oxygen could be mixed with 2001 atoms of hydro- 
gen, after combination had taken place one atom of hydrogen 
would be left uncombined. In this way the theory explains 
very satisfactorily the law of constant proportion. It explains 
equally well the law of combining weights, for these combining 
weights must be directly connected with the relative weights 
of the atoms of the elements. 

In accordance with the atomic theory we may select some ele- 
ment as our unit for atomic weights, and by determining the 
amounts of other elements which combine with a given weight of 
this element and the number of atoms of each element in the 
compounds formed, we can determine the weights of the atoms 
of the other elements as compared with the weight of an atom 
of the element taken asa unit. Thus if we take hydrogen as our 
unit and find that hydrochloric acid contains one part of hydro- 
gen to 35.5 parts of chlorine, and can show, further, that a mole- 
cule of hydrochloric acid contains one atom of chlorine and one 
atom of hydrogen (p. 92), the atom of chlorine must be 35.5 
times as heavy as the atom of hydrogen and we say that the 
atomic weight of chlorine is 35.5. Or if we find that water con- 
tains 8 parts of oxygen for one of hydrogen! and a molecule of 
water contains one atom of oxygen and two atoms of hydrogen, 
the atom of oxygen must be 16 times as heavy as the atom of 
hydrogen and we say that the atomic weight of oxygen is 16. 

The atomic theory, which could be considered as scarcely 
more than a doubtful hypothesis when it was first proposed by 
Dalton, became the central, guiding principle in the development 
of the science of chemistry during the nineteenth century ; and 
evidence in its favor has been accumulated from very many dif- 
ferent and independent directions, so that, now, the actual 
existence of atoms and molecules can scarcely be doubted. 

1 The exact composition of water will be considered later. 


16 A TEXTBOOK OF CHEMISTRY 


We even have a half dozen different ways of estimating the 
actual weight of an atom and the estimates agree fairly well. 
These estimates give the number of molecules in a cubic centimeter 
of air under standard conditions as about 2.71 * 10 or nearly 
thirty million million millions (Millikan. See also Rutherford, 
Presidential Address before Section A of the British Association 
at the Winnipeg meeting). Sir William Thomson (known later 
as Lord Kelvin) once used the illustration that if a drop of water 
could be magnified to the size of the earth the molecules would 
be larger than small shot and smaller than cricket balls. This 
is something the same sort of an estimate as if we were to say that 
a certain animal is the size of adog. Our knowledge of the space 
filled by a molecule is now much more accurate. 

Selection of Atomic Weights. In the series of compounds 
used to illustrate the law of combining weights, the combining 
weights of oxygen are 8, 16 and 48. If the table were extended, 
the values 4 and 32 might be found in other compounds, and 
almost any multiple of 8. It is evident that if we start with 
hydrogen and give it an atomic weight of 1 (see, however, p. 72), 
only one of these various combining weights can be the true 
atomic weight of oxygen. Since the atoms and molecules are 
so small as to be beyond the possibility of direct observation, it 
seemed for a long time impossible to select the true atomic weight 
from among the various possible combining weights: Dalton 
thought it most natural to suppose that the molecule of water 
contains one atom of hydrogen and one atom of oxygen and on 
this basis the atomic weight of oxygen would be 8 instead of 16. 
The reasons for considering that the true atomic weight of oxygen 
is 16 and the methods used in selecting what are believed to be 
true atomic weights will be considered later (p. 92). 

Formulas. The atomic weights selected for the elements used 
to illustrate the law of combining weights are:! H = 1, 
Ca 63.695 ='322 1G) = 8bb he = 56 FN ae ewe 

1'These values are rounded off. The accepted values ars: 


H’= "1.008, O'= 16.00, Cus 68.578 = 32.0640 C1 = 35457 epee 
55.84, Na = 22.997. 


INTRODUCTION LF 


express the composition of water, cuprous oxide and cupric 
sulfide in such a manner as to avoid the use of fractions of atomic 
weights, the ratios for these compounds become : 


Water, He Ose e216 
Cuprous oxide, .O; Cui 916: 127.2 
Cupric sulfide, Cu:S 63.05.02 


In accordance with the atomic theory it follows from these 
ratios that a molecule of water contains two atoms of hydrogen 
for each atom of oxygen, that a molecule of cuprous oxide con- 
tains two atoms of copper for one of oxygen and cupric sulfide 
contains the same number of atoms of sulfur as of copper in its 
molecule. It has been found very convenient to express these 
relations by using the symbol of each element to stand for one 
atom of the element and so to write formulas for compounds, 
using numerical subscripts to designate the number of atoms of 
each element contained in a molecule of the compound. The 
formulas for the compounds are: H2O, Cu,O, CuS. Since a 
formula is always based on the proportion by weight of each 
element contained in the compound, it tells us not only how 
many atoms of each element are contained in a molecule of the 
compound, but it also tells us the exact composition of the com- 
pound by weight. Thus the formula H.SQOu,, for sulfuric acid, 
means that a molecule of sulfuric acid contains two atoms of 
hydrogen, one atom of sulfur, and four atoms of oxygen; but 
it also means that the acid is composed of 2 parts by weight 
of hydrogen, 32 parts of sulfur and 64 parts of oxygen. 

Strictly speaking, the formulas given should be written H.O,, 
Cu20;, Cu;S; and H.S,;O,, but by common consent the subscript 
1 is always understood when no subscript is given. 

What are the formulas of the other compounds mentioned on 
p. 14? ) 

Composition of Pure Substances. From what has been stated 
“we may derive a fifth characteristic of a pure substance. The 
composition of a. pure substance can always be expressed by 
exact multiples of the atomic weights of the elements compos- 


18 A TEXTBOOK OF CHEMISTRY 


ing it. This may be considered as still another way of stating 
the law of combining weights. The law has been tested by the 
analysis of thousands of compounds, and, like the law of constant 
proportion, it is one of the perfect laws from which no deviation 
has been discovered. 

Study of Chemistry. To obtain a knowledge of the elements 
of chemistry it is necessary to become acquainted with a large 
number of facts about the substances with which the science 
deals, but it is still more important to understand the relations 
connecting these facts with each other and the fundamental 
laws and theories by which the facts are grouped together and 
explained. Success in the study depends especially on the abil- 
ity to learn new facts in their relation to those which have al- 
ready been acquired and on the cultivation of a logical as dis- 
tinguished from an arbitrary memory. Formulas, especially, 
should be derived, whenever possible, from the formulas of other 
compounds of the same elements, and not learned individually, 
except in the earliest portion of the study. Chemical equations 
should be written on the basis of a knowledge of the reacting 
substances and of the products of the reaction, and should never 
be learned by brute memory. 

In the systematic treatment of the subject the more common 
elements will be considered first and under each element the 
compounds of that element with each of those previously studied 
will be mentioned so far as this is desirable. See p. 132. 


CHAPTER II 


OXYGEN 
SymsBot, O. Atomic WeicuHT, 16. 


Occurrence. Oxygen is the most abundant and one of the 
most important of all the elements. It forms about one fifth 
of the volume of the air, eight ninths of the weight of water and 
nearly one half the weight of the mineral substances which com- 
pose the crust of the earth. Oxygen is found in all living bodies 
and is a constituent of a larger number of compounds than any 
other element except carbon. 

Preparation. 1. When metallic mercury is heated at the 
right temperature in contact with the air, it is slowly converted 





FG) G 
WORN 


c i in 
SO 
sii 











ae 
EE (yy 











HH 
if 
| 
| 


i{ 


ee 


EN 
i) 
Hy 

/ 














i 


i 
MULL 


aif 
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into a bright red compound called oxide of mercury. The French 

chemist, Lavoisier, carried out the experiment in the apparatus 

shown in the figure, and proved that after the mercury had been 

heated several weeks the air no longer decreased in volume and 

he concluded that this was because the oxygen of the air had 
19 


20 A TEXTBOOK OF CHEMISTRY 


all been removed by combination with the mercury. He then 
collected the oxide of mercury and heated it to a higher tempera- 
ture till it was all decomposed into mercury and gaseous oxygen. 
The volume of oxygen was the same as the decrease in volume 
of the air during the heating in contact with the mercury. 

The quantitative relation between the mercury, oxygen and 
oxide of mercury may be very briefly expressed by means of the 
symbols for the elements, as follows: 


Hg ae O. = HgO 
Mercury _ Oxygen Oxide of Mercury 
HgO 4 Hens O 
Oxide of Mercury Mercury Oxygen 


Since the atomic weight of mercury is 200 and the atomic 
weight of oxygen is 16, the first equation means that 200 parts 
by weight of mercury combine with 16 parts by weight of oxy- 
gen to form 216 parts of oxide of mercury ; and the second equa- 
tion means that 216 parts of oxide of mercury decompose into 
200 parts of mercury and 16 parts of oxygen. It will be noticed 
that the symbols of two elements placed side by side represent 
a compound, while a symbol by itself represents a free 
element. 

2. The portion of air which was not absorbed by the mercury 
was chiefly nitrogen and formed about four fifths of its volume. 
Liquid nitrogen boils at — 194°, while liquid oxygen boils at 
— 182.5°. If liquid air is allowed to boil, the nitrogen goes off, 
chiefly, at first, and the gas which comes off toward the end is 
nearly pure oxygen. In this way the oxygen and nitrogen may 
be separated very much as alcohol and water are separated by 
distillation. Oxygen prepared in this manner is compressed into 
strong steel cylinders for medicinal and other uses. 

3. When potassium chlorate is heated, it melts and begins 
to decompose slowly into potassium chloride and oxygen. 
Potassium chlorate has the composition represented by the 


formula KCIOs3. 


OXYGEN 21 


Potassium, K = 39.1 parts or 31.90 per cent 
Chlorine, Cl = 35.46 parts or 28.93 per cent 
Oxygen, 30 = 48. partsor 39.17 per cent 

Total 122.56 parts or 100. per cent 





The decomposition may be represented by the equation: 
KC1lOs = VAP K Gi wii e 2310 


Potassium Potassium Oxygen 
Chlorate Chloride 


If the potassium chlorate is mixed with one fourth of its 
weight of finely powdered manganese dioxide, MnO., the de- 
composition will begin at a much lower temperature and pro- 
ceed more rapidly than when the potassium chlorate is heated 
alone. If the residue in the retort, after the decomposition is 
complete, is treated with water, the potassium chloride will 
dissolve, while the manganese dioxide will remain undissolved 
and may be readily separated from the solution by filtration. 
If the manganese dioxide is examined, it will be found that it 
has not changed in composition or amount. We may, there- 
fore, write the equation : 


KCIO; + MnO, = KCl + 30 + MnO, 
Potassium Manganese Potassium Oxygen Manganese 
Chlorate Dioxide Chloride Dioxide 


4. When fused sodium peroxide containing a very little 
copper oxide! is dissolved in water, it gives sodium hydroxide 
and oxygen : 


Na,O, 4. H,O = z NaOH -|- O 
Sodium Water Sodium Oxygen 
Peroxide Hydroxide? 


As the copper oxide is left unchanged at the end and as the 
reaction will take place, though more slowly, in its absence, we 
may omit it in writing the equation. ; 


1 The substance is known commercially as ‘‘ oxone.”’ 

2 The student should notice the connection between the name and 
the composition of sodium hydroxide. Many similar compounds 
containing oxygen and hydrogen are called hydroxides. 


22 A TEXTBOOK OF CHEMISTRY 


Collection and Storage of Gases. Gases which are only 
slightly soluble in water and which are not required in a high 
state of purity are conveniently collected and 
stored in a gasometer of the form shown in 
Fig. 3. If the stopcocks A and B are opened 
and the cap C’ screwed on, water placed in the 
cup above will run into the body of the gas- 
ometer till it is filled with water and all of 
the air is expelled. Now on closing the stop- 
cocks the gasometer will still remain filled with 
water after removing the cap at C. By means 
of a tube inserted through C gas may be in- 
troduced and will fill the gasometer, displacing 
the water, which will flow out of C by the 
side of the tube delivering the gas. After 
filling the gasometer and replacing the cap at 
C, on opening the stopcock A, water will run 
into the body of the gasometer and the gas 

Fig. 3 may be drawn off as desired through B. 
Properties of Oxygen. Oxygen is a colorless, 
odorless and tasteless gas.1 The weight of one liter at 0° 
and under a pressure of 760 millimeters of mercury (about 
the average atmospheric pressure at sea level) is 1.429 grams.? 
Under the same conditions of temperature and pressure it is 
about one tenth heavier than the same volume of air. 

The most striking property of oxygen is the vigor with 
which it supports combustion. All substances which burn in 
air burn much more rapidly and vigorously in oxygen. A splin- 
ter of wood having a live coal on the end will burst into flame, 





1These statements refer, of course, to the ordinary form of the 
element. Ozone, which is another form of oxygen, is colored and has 
a strong odor (p. 98). 

2 At 45° latitude. It is slightly less at lower latitudes because the 
pressure of 760 mm. of mercury is less and the density of the gas is 
less. At the latitude of New York the weight of one liter of oxygen 
is 1.42845 grams. If, however, the reading of the barometer is cor- 
rected for latitude and altitude, the weight of one liter of oxygen is 
1.429 grams at any place. 


OXYGEN 23 


if thrust into the gas. A piece of charcoal, barely ignited, will 
glow intensely and be surrounded by a pale blue flame, scarcely 
visible in the intense light of the glowing mass. 
The product of the combustion is carbon di- 
oxide, a colorless gas. Sulfur burns with a 
brilliant blue flame, giving sulfur dioxide, also 
a colorless gas. Phosphorus burns with an 
intense white light, giving a white, solid com- 
pound, phosphorus pentoxide. A coil of iron 
wire or a steel watch spring to which is at- 
tached a string that has been dipped in melted 
paraffin, may be set on fire and will burn in 
oxygen (Fig. 4), throwing off brilliant sparks 
and forming white-hot, molten globules of the 
magnetic oxide of iron, which will drop off 
from time to time. 





Fig. 4 


The equations which represent the quantitative relations in 
these experiments are: 


Carbon Oxygen Carbon Dioxide —- 
5 ee = SO, 
Sulfur Oxygen Sulfur Dioxide 
Die es sponte, yy tes P.O; 
Phosphorus Oxygen Phosphorus Pentoxide 
3 Fe + 40 = Fe;,04 
Tron Oxygen Magnetic Oxide of Iron 


Oxygen and Acid Properties. If sulfur dioxide or phosphorus 
pentoxide is dissolved in water, the solution obtained will have 
a sour taste, and acid properties. Many other compounds of 
nonmetallic elements with oxygen combine with water in a 
similar manner to form acids, and it is because of this that the 
name oxygen, meaning “ acid former,’ was first given to the 
element. When the name was given, it was supposed that all 
acids contain oxygen, but it was discovered later that this is 
not the case. 


24 A TEXTBOOK OF CHEMISTRY 


~ 


Combustion. Effect of Concentration on a Chemical Reac- 
tion. The similarity between ordinary combustion and the 
burning of substances in oxygen is apparent. A more careful 
study of the matter has shown that exactly the same compounds 
are formed when charcoal, sulfur or phosphorus burn in the air 
as are formed when they burn in oxygen, and even in the case of 
iron, the magnetic oxide formed by burning the steel watch 
spring has just the same composition as that of the scale formed 
when white-hot iron is exposed to the air. The burning of iron 
in air can also be shown by sprinkling 
fine iron filings through a flame. 
About four fifths of the air by volume 
is nitrogen. This does not combine with the 
burning substances, and by its presence it 
moderates the action of the oxygen, partly 
by diluting it, partly because it must be 
heated to the same temperature as the other 
substances, and this absorbs a large part of 
the heat of the reaction and so lowers the 
temperature to which the burning substance 
is heated. These facts illustrate two prin- 
ciples of almost universal application in 
chemistry: first, that the speed of a chemi- 
cal reaction is increased by increasing the 
concentration of one of the reacting sub- 
stances, here the concentration of the oxygen; and, second, 
that the speed of a reaction is affected by the temperature and 
is greater at high temperatures than at low ones. We shall find 
later that the first of these principles can be stated in the form 
of an accurate, quantitative law, but the phenomena of ordi- 
nary combustion are not well suited for a quantitative study 
of this kind. 

Kindling Temperature. If we place on an iron plate smail 
pieces of phosphorus, sulfur and charcoal, it will be found oa 
warming the plate that phosphorus takes fire at a quite low 
temperature, the sulfur at a moderate heat, while the charcoal 





CALORIMETER 20 


will not burn till the plate is nearly red-hot. The temperature 
at which combination of a substance with the oxygen of the air 
is sufficiently rapid so that it takes fire is called the kindling 
temperature. ‘The temperature rises rapidly from the heat of the 
reaction as soon as the substance is kindled. Well-known appli- 
cations of the gradations of kindling temperature are the old- 
fashioned sulfur match and the methods commonly used in kin- 
dling a fire. Kindling temperature is not a satisfactory measure 
of the affinity of a substance for oxygen nor is it closely con- 
nected with the heat generated on combination with oxygen. 
Heat of Combustion. Calorimeter. When substances burn, 
a part of the chemical energy of the burning substance and of 
the chemical energy of the oxygen is converted into heat. The 
amount of energy transformed into 
heat when one gram of the substance 
burns, or, for most scientific pur- 
poses, the energy obtained from one 
gram atom or gram molecule of the 
substance, is called its heat of com- 
bustion. By gram atom is meant as 
many grams of the substance as 
there are units in the atomic weight 
— as 12 grams of carbon or 31 grams 
of phosphorus. By gram molecule is 
meant as many grams of the sub- 
stance as there are units in the mo- 
lecular weight, as 44 grams of carbon 
dioxide (12 grams of carbon + 32 
grams of oxygen). 
The heat of combustion is de- 
termined in an instrument called a 
calorimeter, the most accurate form 
being known as a bomb-calorimeter 
because the combustion is carried 
out in an atmosphere of compressed oxygen in a strong, tightly 
closed bomb, which is immersed in water. ‘The substance is 





Fig. 6 


26 A TEXTBOOK OF CHEMISTRY 


placed in a small crucible within the bomb and is ignited by 
means of a fine iron wire, which is heated for a moment by an 
electric current. The weight of the substance, the weight of 
water surrounding the bomb and the temperature of the water 
before and after the substance is burned are accurately de- 
termined. There are, of course, many other details about the 
apparatus and manipulation which need not be described here. 
(See Atwater, J. Am. Chem. Soc. 25, 659, and Richards and 
Jesse, ibid. 32, 268.) 

The amount of heat required to raise the temperature of one 
gram of water one degree at 15° is called a calorie (see p. 33). 
This is often called the small calorie and designated by the 
abbreviation cal. to distinguish it from the large Calorie, which 
is the amount of heat required to raise the temperature of a 
kilogram of water one degree and which is designated by the ab- 
breviation Cal. 

* The corresponding unit of the English system is the British 
Thermal Unit (B. T. U.), the heat required to raise the tempera- 
ture of a pound of water 1° F. Since 1 Kg. = 2.204 lb. and 1° F. 
= 3 of 1°C., 1 Cal. = 3.968 B. T. U. Heat of combustion ex- 
pressed in calories per kilogram, however, is reduced to British 
Thermal Units per pound by multiplying by %. From the re- 
sults of determinations with the calorimeter it is possible to cal- 
culate how many grams of water can be raised one degree in tem- 
perature by burning one gram or one gram atom of a substance 
and this will be the heat of combustion of the substance in 
calories. 

The heats of combustion for the substances mentioned in this 
chapter are: 


1Since the amount of heat required to raise the temperature 
of one gram of water one degree varies slightly with the tempera- 
ture, it is necessary in accurate work to define the temperature 
at which the experiment is performed. A mean temperature of 
15° is most often used. On account of the variability of the 
calorie it has been proposed to use the joule as a unit. The small 
calorie is equal to 4.182 joules at 15°. See G. N. Lewis, Journal of 
the American Chemical Society, 35, 4 (1913). 


CHEMICAL ENERGY Zi. 


For one gram of carbon (charcoal), 8080 calories 


For one gram of sulfur (rhombic), 2190 calories 
For one gram of phosphorus (yellow), 5970 calories 
For one gram of iron (to Fe30,), 1616 calories 


For one gram of mercury (to HgO), —_‘105 calories 


For one gram atom of carbon (charcoal), 96,960 calories 
For one gram atom of sulfur (rhombic), 70,180 calories 
For one gram atom of phosphorus (yellow), 185,000 calories 
For one gram atom of iron (to Fe3Qx), 90,200 calories 


For one gram atom of mercury (to HgO), —21,000 calories 


The Nature of Chemical Energy. The amount of energy 
liberated by a burning substance is very large. If it were 
possible completely to transform the energy liberated by burning 
a pound of good coal into mechanical energy, it would lift a ton 
weight over 4000 feet. It is a very good engine which will con- 
vert ten per cent of the energy of the coal burned under its boiler 
into useful work, but, in spite of the fact that more than ninety 
per cent of the energy of the coal is dissipated and lost, the total 
energy is so enormous that the steam engine is commercially 
economical. 

A very natural question which arises here is; What is the 
source of the energy which suddenly appears as heat when par- 
ticles of carbon and oxygen unite to form carbon dioxide? Is 
there some motion within the particles of oxygen and carbon 
which is transformed into heat when they unite? or do the 
particles collide when their mutual attraction brings them to- 
gether, somewhat as a meteor collides with the earth? For 
these questions there are, at present, no answers, and specula- 
tions about them are of very little value till some one can dis- 
cover some sort of experimental evidence bearing upon them. 
It is well, however, to recognize how imperfect and fragmentary 
our knowledge is and that there are hundreds of questions like 
these for which we have. no answer. It is also well, at times, 
to ask such questions and consider whether there is any tangible 
method of attacking the problem, for, while the explanation seems 


28 A TEXTBOOK OF CHEMISTRY 


beyond our grasp, at present, many similar problems which 
would have seemed beyond the possibility of a solution one hun- 
dred years ago have been solved. 

Catalysis. It has been pointed out that when manganese 
dioxide is mixed with potassium chlorate the latter decomposes 
at a lower temperature or more rapidly than when the chlorate 
is heated by itself, but that the manganese dioxide is left un- 
changed in the end. When a substance acts in this manner by 
its mere presence, causing a reaction or decomposition to take 
place at a lower temperature or more rapidly, it is called a 
catalytic agent and the process is called catalysis. 

* A study of this particular case makes it seem probable that 
the oxygen is at first transferred from the potassium chlorate to 
the manganese dioxide and that the compound of the manganese 
dioxide with the oxygen decomposes at a lower temperature 
than the potassium chlorate. Such an explanation seems, at 
first, paradoxical, for it seems to imply that manganese dioxide 
has a greater affinity for oxygen than potassium and chlorine 
have and so can take the oxygen away from the potassium 
chlorate, while in the resulting compound the affinity of the 
manganese dioxide for oxygen seems to be less than that of 
potassium and chlorine for oxygen, because the decomposition 
of the manganese compound occurs at a lower temperature 
than that required for the decomposition of potassium chlorate. 

A partial explanation of this seeming paradox is found in the 
fact chat very many substances act upon each other chemically 
at a much lower temperature than that at which either decom- 
poses into its elements. It is also true that the stability of a 
compound is not an accurate measure of the affinity between the 
elements of which it is composed. The affinity between the 
elements of a compound which decomposes at 200° is not neces- 
sarily less than that between the elements of a compound which 
decomposes at 400°. Still further, it is by no means always true 
that when an element is transferred from one compound to 
another its affinity for the element with which it combines is 
greater than that for the element which it leaves. These ques- 


CHEMICAL AFFINITY. NOMENCLATURE 29 


tions will be considered further later ; but it is well, at the outset, 
to avoid certain misconceptions which are very liable to arise 
because the facts of chemistry are often so very different from 
what our first and most natural idea of chemical affinity would 
lead us to expect. 

Chemical Affinity. The term affinity has been used in the pre- 
ceding paragraph and seems to call for some definition. The 
word is generally used in a rather indefinite way to designate 
that attraction between elements which causes them to unite to 
form compounds. 

* Our knowledge of the real nature of chemical affinity is very 
vague and indefinite. The theory that atoms consist of nuclei 
carrying a positive charge of electricity surrounded by electrons, 
which are atoms of negative electricity, must be considered as 
quite well established. It seems certain that the electrons are 
intimately connected with the forces which hold the atoms 
together in compounds and several hypotheses have been pro- 
posed to account for the arrangement of the electrons. None 
of these gives reasons which are at all definite for the varying 
attractions between different atoms and there has been no 
attempt to explain, accurately, the mechanism by means of 
which heat is generated when elements combine. 

It is possible, however, ‘to learn a great deal about the condi- 
tions under which chemical affinity acts. Thus it seems always, 
in accordance with the laws of constant proportion and com- 
bining weights, to be exerted between the atoms and, for this 
reason, between definite quantities of the elements. By study- 
ing equilibria in chemical reactions it is, also, sometimes possible 
to give affinity a mathematical definition and measure it 
quantitatively. 

Nomenclature. The compounds of oxygen have been called, 
in this chapter, oxides. This is an application of a system of 
naming substances which is used for all compounds consisting 
of two elements. As compounds of oxygen are called oxides, 
compounds of sulfur are called sulfides, compounds of chlorine, 
chlorides, etc. In order to give more definite names, prefixes 


30 A TEXTBOOK OF CHEMISTRY 


derived from the Greek numerals are used. CO is called carbon 
monoxide; COs, carbon dioxide; HgCle, mercury dichloride ; 
P.O3, phosphorus trioxide; SOs, sulfur trioxide; CCl4, carbon 
tetrachloride; P2O;, phosphorus pentoxide. In addition to 
these names, which tell how many atomic weights of the ele- 
ment are contained in the molecular weight of the substance, 
the prefix per is used to name compounds containing more oxy- 
gen than some other oxide of the same element. Thus sodium 
peroxide, Na2O2, contains more oxygen than the other oxide of 
sodium, Na2O, the prefix per meaning more or beyond. 

Still another method of naming oxides is to add the suffixes 
-ous and -ic to the name of the metal or other element which is 
combined with oxygen. The ending -ic is used for the compound 
containing the larger proportion of oxygen. Thus Hg20 is called 
mercurous oxide; HgO, mercuric oxide; FeO, ferrous oxide; 
Fe.03, ferric oxide. 

The choice among these three methods of naming oxides and 
other compounds is more or less arbitrary and conventional. 


CHAPTER III 
LAWS OF GASES 


Units or LenetH, WericHT, VoLuME, TEMPERATURE, TIME AND 
ENERGY. 


Unit of Length. Meter. ‘The meter was originally intended 
to be one ten-millionth of the distance from the equator to the 
pole of the earth, measured on the surface. The measurements 
by means of which the first meter was prepared were inaccurate, 
however, and the real meter is the distance, measured at the 
freezing point of water, between two marks on a bar of platinum- 
iridium kept at the International Bureau of Weights and Meas- 
ures at Sévres, France. The meter is divided into tenths, hun- 
dredths and thousandths, called decimeters, centimeters and 
millimeters. Its most common multiple is the kilometer, 
1000 meters. 

Unit of Weight. Gram. The gram was intended to be the 
weight of one cubic centimeter of water at its maximum density, 
4° centigrade. Here, again, the first measurements were not 
quite accurate and the real kilogram (1000 grams) is the weight, 
in a vacuum, of a mass of platinum-iridium kept at the Inter- 
national Bureau. ‘The most common division of the gram is one 
milligram, the thousandth of a gram. 

Unit of Volume. Liter. The liter was intended to have a 
volume of one cubic decimeter. Because volumes can be most 
accurately compared by weighing the water which fills them, 
the real liter is the volume occupied by one kilogram of water, 
weighed in a vacuum at 4° C. The cubic centimeter is defined, 
conventionally, as one one-thousandth of a liter.1 


1The actual weight of water contained in a cube whose edge is 
one centimeter is 0.999982 g. according to the best measurements. 
Because the edge of a cube of water weighing one gram is not ex- 


31 


32 A TEXTBOOK OF CHEMISTRY 


Units of Time. The units of time used in chemistry are the 
second, minute, hour, day and year. These are all fixed by 
means of astronomical observations with the aid of accurate 
clocks. 

Unit of Temperature. ‘The freezing point of water under at- 
mospheric pressure has been selected as zero for the ordinary 
centigrade scale of temperature, and the boiling point of water 
under atmospheric pressure as 100°. For the International scale, 
the interval between the two points is divided into one hundred 
equal parts by measuring the increase in pressure of hydrogen 
gas, at constant volume, the initial pressure being that of a 
column of mercury one meter high at 0°. Absolute tempera- 
tures will be considered later. 

It is well to notice that the unit of temperature is a unit of 
intensity and not of quantity. In this respect it corresponds 
to the height to which a weight is raised in mechanical energy 
or to the volt in electricity. 

Units of Energy. Kilogram-meter; Erg. The simplest unit 
of energy is the kilogram-meter, the energy required to lift one 
kilogram to a height of one meter. Since the force of gravity 
varies with the latitude and altitude, another unit, which is 
independent of these, is often used. ‘This is the erg, and is twice 
the energy of one gram ! moving with a velocity of one centimeter 
a second. Or it may be defined as twice the energy required to 
impart to one gram a velocity of one centimeter per second or 
to increase its velocity by one centimeter per second. One joule 
is 10,000,000 ergs. 


actly one centimeter in length, some persons prefer to call the con- 
ventional cubic centimeter a milliliter. The suggestion has not, 
however, been generally accepted. 


1 This is more often stated as the mass of one gram, but since 
weights are always accurately determined by the balance, one gram 
determined by weighing is just as ‘‘absolute”’ a quantity of matter 
as the mass of one gram. For the same reason chemists are justi- 
fied in speaking of atomic weights instead of atomic masses. It is 
well to remember, however, that weight is in its accurate, scientific 
use, a@ measure of a force and not a measure of a quantity of matter. 
Me ne tN everyday use of the word we use it for a quantity 
of matter. . 


UNITS 533) 


Centimeter-gram-second System. Absolute units. Since in 
accordance with the law of conservation of energy every form of 
energy bears an exact, quantitative relation to every other, any 
quantity of energy which we can measure may be expressed in 
terms of the velocity of a moving mass. The units necessary 
for such a purpose are a unit of length, a unit of mass and a 
unit of time. Physicists have agreed upon the centimeter, 
gram and second as fundamental units and have developed a 
system of “‘ absolute units’ in which all forms of energy are 
measured by reference to these. These units are called absolute 
because they are independent of the force of gravity. 

Units of Mechanical Energy. The absolute unit for mechan- 
ical energy is the erg, which has been defined above. The most 
common unit used by engineers is the kilogram-meter (or the 
foot-pound in the English system). At 45° latitude and sea 
level the kilogram-meter is 98,066,700 ergs. 

Unit of Power. Power is the rate of production of energy. 
One horse power is 4600 kilogram-meters or 33,000 foot-pounds 
per minute. 

Units of Heat. ‘The calorie is the heat required to raise the 
temperature of a gram of water one degree (p. 26). It varies 
slightly with the temperature and for accurate work the tem- 
perature must be specified — usually a temperature at 15°! is 
taken as the standard. As an absolute unit, independent of 
the temperature of the water the joule has been suggested. 
One calorie at 15° is equal to 4.187 joules. 

Electrical Units. The primary electrical units are the volt, 
ohm and ampere. These are so related that an electromotive 


1Tf the calorie at 15° is taken as one, the values of the calorie at 
other temperatures are as follows: 


10° 1.0016 
15° 1.0000 
20° 0.9991 
25° 0.9988 
30° 0.9989 


These are the mean of the values of Georg Ianke, Ann. Tables of 
Physical Constants for 1910, and of Bausfield, Phil. Trans. 271, A, 
199 (1911). 


34. A TEXTBOOK OF CHEMISTRY 


force (E. M. F.) of one volt acting through a resistance of one ~ 
ohm gives a current of one ampere or: 


POE OS 


A ee pas 
ate ee Herrin 


The unit of electrical power is the watt, a current of one ampere 
flowing under a difference of potential of one volt. It is equiva- 
lent to 10,000,000 ergs or one joule per second. The kilowatt 
is, of course, 1000 watts and is the most common measure for 
electrical service in lighting, heating, running of motors and the 
like. The watt is one of the “ absolute ” units. 

An electrical horse power is 746 watts, and is equivalent, of 
course, to 33,000 foot pounds per minute. 

Chemical Energy. By chemical energy is meant the energy 
which appears as mechanical energy, heat, light, sound or elec- 
tricity when two or more elements unite,! or when an element is 
changed from one form to another, as ozone to oxygen. It is 
usually expressed in terms of heat units or electrical units. It 
must always refer to some definite chemical action which takes 
place and can never refer to the total energy contained in an ele- 
ment or compound, as we have no means of measuring this. 

Effect of Pressure on a Gas. Law of Boyle. When the pres- 
sure applied to a gas is doubled, the volume is reduced to one 
half ; or when the pressure is reduced to one half, the volume be- 
comes twice as great. Another method of stating this property 
of gases in a general way is to say that the volume of a gas varies 
inversely as the pressure. Or, mathematically : 


Vat PR ePs oro Rese Ps onerane 


where V and V’ are two volumes of the same quantity of gas and 
P and P’ are the corresponding pressures. This is known as 
Boyle’s law. It is not an accurate law, as the law of constant 


1 Heat is absorbed when some substances unite, and in such cases 
the energy of the compound is considered as negative or less than 
cothing in comparison with that of the elements from which it is 
formed, but the idea that energy ean be really negative seems 
absurd. 


LAWS OF GASES oo 


prvportion is, but is sufficiently accurate for use in all ordinary 
cases. 

* The extent of the deviation from the law for several gases is 
shown in the following table: 


TABLE 


Volumes filled at 0° by two liters of each gas when the pressure is 
increased from one atmosphere to two atmospheres. 


Hydrogen - 1.0006 liters 
Nitrogen 0.9996 liters 
Carbon monoxide 0.9995 liters 
Oxygen 0.9991 liters 
Nitric oxide 0.9989 liters 
Carbon dioxide 0.9931 liters 
Nitrous oxide 0.9924 liters 
Hydrochloric acid 0.9919 liters 
Ammonia 0.9845 liters 
Sulfur dioxide 0.9739 liters 


Those gases which are liquefied most easily depart farthest 
from the law, and all gases except hydrogen and helium are com- 
pressed more than they should be’under moderate pressures. 
For a pressure of many atmospheres a point is reached where all 
gases which do not liquefy are compressed less than they should 
be in accordance with the law. According to the kinetic theory 
(p. 58) the greater compressibility under moderate pressure 
is caused by the attraction of the molecules for each other — 
the same forces which cause the gas to liquefy at low tempera- 
tures or under pressure. The point of too little compressibility is 
reached when the molecules are brought so close together that the 
molecules themselves fill a considerable fraction of the total space. 

The law may be easily illustrated by taking a gas measuring 
tube, graduated in cubic centimeters, filling it partly full of 
mercury and immersing the mouth in a deep, narrow jar con- 
taining mercury. It is evident that if the tube is raised or 
lowered till the top of the mercury within the tube is exactly 
level with the surface of the mercury on the outside, the pressure 


36 A TEXTBOOK OF CHEMISTRY 


of the gas within the tube will be the same as that shown by 
a barometer in the same room.! If, now, the tube is 
raised, the volume of the gas will be seen to in- 
crease, and for any given position the pressure of 
the gas must be equal to the reading of the barome- 
ter less the height of the mercury in the tube above 
that in the jar. By reading the volumes in two 
different positions of the tube and determining the 
corresponding pressures the data for a verification of 
the law may be easily obtained. 

For practical uses it is convenient to select some 
standard pressure to which the volume of a gas 
may be referred. The pressure universally used by 
chemists for this purpose is the pressure of a column 
of mercury 760 mm. high at 0°, at 45° latitude and 
at sea level. This is, approximately, the average 
pressure of the air at sea level and is called a pres- 

Fig.7 sure of one atmosphere. Other pressures are most 

easily determined by measuring, directly or indi- 

rectly, the height of the column of mercury which will balance 
the elastic pressure of the gas. 





* Corrections for Readings of the Barometer. — In accurate work, 
when the barometer is read at some other temperature than 0° a cor- 
rection must be subtracted, owing to the fact that the column of 
mercury is lighter as the metal expands with rise of temperature. ‘The 
correction in millimeters at temperatures from 5° to 35° is: 














TEMPERATURE CORRECTION FOR BAROMETER | CORRECTION FOR BAROMETER 
DEGREES WITH GLAss SCALE witTH Brass ScAaLE 

5 0.7 0.6 
10 1.3 122 
1D 2.0 1.9 
20 2.6 2.5 
25 3.3 3.1 
30 4.0 SWE 
30 4.7 4.3 








1 For the sake of simplicity, the lowering of the meniscus of the 
mercury in the tube by eapillary action is disregarded. 


LAWS OF GASES 37 


If the pressure is less than 760 mm. the correction will be less in pro- 
portion. Thus at 730 mm. the correction for a glass scale is 2.5 mm. at 
20° instead of 2.6 mm. 

The corrections for latitude and altitude are usually less important. 

Corrections of barometer for latitude, to be added for latitudes less 
than 45° or subtracted for latitudes greater than 45°: 

















LATITUDE CORRECTION LATITUDE 
Ua 1.97 90° 
iy 1.94 85° 

10° 1.85 80° 
ARs be 1.70 (Gy 
an 1.51 70° 
25° ieee 65° 
30° 0.98 60° 
35° 0.67 55° 
40° 0.34 50° 
45° 0.00 45° 








Correction for altitude, to be added. 

















ALTITUDE CORRECTION BAROMETER READING 
300 meters 0.04 720 
600 meters 0.08 700 
900 meters 0.12 680 
1200 meters 0.16 660 
1500 meters 0.19 640 
2000 meters 0.24 630 























If the corrections for latitude and altitude are applied to the barometer 
readings, the weight of one liter of the gas at 45° latitude may be properly 
used in calculating the weight of a quantity of gas measured at any other 
latitude or altitude. 


A problem which often presents itself in dealing with gases 
is the calculation of the volume which a quantity of gas, that 
has been measured at some other pressure than that of one at- 
mosphere, would assume if it were brought to atmospheric pres- 


38 A TEXTBOOK OF CHEMISTRY 


sure. Such problems are most easily solved by putting the 
mathematical expression given above into the following form : 


VP! ya 
iB 760 


y= Or Veatiou nit, sre 





The student is advised most earnestly that this formula should 
not be committed to memory. Instead of this it should only 
be remembered that when the volume at one pressure is given and 
the volume at another pressure is desired, the first volume is to 
be multiplied by a fraction in which one pressure is the numera- 
tor and the other pressure the denominator. A consideration of 
the fact that an increase in pressure will cause a decrease in the 
volume will at once indicate which pressure is to be taken as the 
numerator of the fraction. The proper method of using the 
formula is emphasized because in the study of chemistry it is 
of the greatest importance to cultivate the ability to reason 
quickly from one point to another and to acquire a knowledge 
of the subject by a rational process rather than by mere 
memory. 

Effect of Temperature on a Gas. Law of Charles. When 
the temperature of a gas is increased one degree while the pres- 
sure remains constant, the volume will increase 545 (or 0.003663) 
of its volume at 6°.! This will be most easily understood with 
the aid of the accompanying diagram, which gives the volume 


1This law, while sufficiently accurate for ordinary calculations, is 
only approximate, the deviations from it being of somewhat the same 
order of magnitude as the deviations from the law of Boyle. The 
coefficients of expansion of some of the more common gases as deter- 
mined by the increase of pressure at constant volume are: 


Air 0.003666 or 1/272.8 Argon 0.003668 or 1/272.6 
Oxygen 0.003674 or 1/272.2 Helium 0.003663 or 1/273.0 
Nitrogen 0.003668 or 1/272.6 Carbon mon- 

Nitric oxide 0.003676 or 1/272.0 oxide 0.003667 or 1/272.7 


Hydrogen 0.003663 or 1/273.0 Carbon dioxide 0.003698 or 1/270.4 
Sulfur dioxide 0.003845 or 1/260.1 


As with the law of Boyle, those gases which are easily liquefied vary 
most from the rate of expansion for a “ perfect”’ gas. 


ABSOLUTE TEMPERATURES 39 


which 273 cubic centimeters of a gas at 0° would assume at 
other temperatures. Only hydrogen or helium would obey 
the law at atmospheric pressure over the range of volumes 
given in the diagram. All other gases are liquid or solid 
at (200% 

Absolute Temperatures. On eens Nee Las VoLUME 
the left side of the diagram is a7 100° 373 Go! 
given a series of numbers which } 
are called absolute temperatures. 

A little examination of the dia- 283° 10° 283 ce. 
gram will show that these tem- 273° 0° 273 6C. 
peratures are based on _ the 
thought that if we could find 
a gas which does not liquefy 


and which continued to con- rh ake bscs: 
tract at the same rate at very 

low temperatures it would dis- 

ie} 

appear at olor If we take 799 900° 73.00, 
this point as the starting point 

for the “absolute” scale of 

temperature it is evident that 0° — 273° 

the freezing point of water Fig. 8 


will be at 273° absolute and 

the boiling point 373°. Any temperature may be readily 
converted to the absolute scale by adding to it, algebraically, 
273). 

The absolute scale of temperature enables us to give a very 
simple statement of the law of Charles, viz.: The volume of a 
gas varies directly as the absolute temperature. ‘This becomes, 
mathematically : 
ere 


(ke 


Vie Or Va 





If, as is customary in dealing with gases, we wish to find the 
volume which a gas, which has been measured at some other 
temperature than 0°, would assume if cooled or warmed to zero, 


AO A TEXTBOOK OF CHEMISTRY 


273 
Te 
be used rationally, not by rote (see p. 38), and may be combined 
with the formula for pressures for practical uses. If the volume 
of a gas is known at one temperature and pressure, its volume at 
some other temperature and pressure may be calculated by multi- 
plying by two fractions one of which involves the two pressures 
and the other the two absolute temperatures. 

Significance of the Absolute Zero. The absolute scale of 
temperature may be treated merely as a mathematical conven- 
ience in dealing with problems of gases and of thermodynamics ; 
but the question naturally arises whether the absolute zero has 
any further, real meaning. Is it, in reality, as the name indi- 
cates, a point of absolute cold at which all phenomena of tempera- 
ture begin and below which it is impossible to go? Many differ- 
ent phenomena seem to indicate that the absolute zero is actual 
and not merely a mathematical fiction. It can be no mere acci- 
dent that hundreds of gases and vapors obey the law of Charles so 
closely ; and the further the study of the physical and chemical 
properties of gases is carried, the more clear does it become that 
the law is intimately connected with some of the most funda- 
mental properties of matter. From the side of experiment, also, 
every recent attempt to reach very low temperatures has pointed 
to — 273° as a point which can never be passed. The lowest 
point thus far reached is that of helium boiling under a very low 
pressure and is estimated as — 272°, or 0.9° absolute. (Kamer- 
lingh Onnes, Chemical Abstracts, 1923, p. 2073.) 

Determination of the Weight of a Liter of aGas. The weight 
of a unit volume of any gas under standard conditions is one 
of its most important properties, not only for the purpose of cal- 
culating the weight of a gas when we know its volume, but be- 
cause of relations between these weights for different gases, upon 
which one of the most important laws of chemistry is based 
(p. 89). The unit volume usually chosen is the liter and the 
standard conditions are a temperature of zero and a pressure of 
760 mm. of mercury. 





the formula may be written, Vo= V’ This formula should 


WEIGHT OF GASES 41 


If a bulb! is connected with a manometer and evacuated by 
means of an air pump, by reading the manometer and tempera- 
ture, we can, if we know the capacity of the bulb, calculate the 
volume which the air remaining in the bulb would fill at 0° and 
760 mm. If we weigh the bulb and then fill it with some gas at 





Fig. 9 


atmospheric pressure (to be determined by reading the barom- 
eter) and weigh it again, the difference between the two 
weights will evidently be the weight of the gas which has entered, 
while the volume of the air which was left in the bulb plus that 
of the gas which has entered can be readily calculated for stand- 
ard conditions as before.?, The difference between this calcu- 
lated volume and the corrected volume of the air which re- 
mained in the bulb will be the volume, under standard con- 
ditions, of the gas which was admitted. From this and the 
weight it is easy to calculate the weight of one liter of the gas 
under standard conditions. 


1A capacity of 125 to 150 ce. is suitable for a lecture or labora- 
tory experiment. The volume may be determined by weighing 
the bulb empty and then filled with water, but the bulb must be 
thoroughly dried by warming it and evacuating it repeatedly before 
it is used for the determination. 

2 This assumes, of course, Dalton’s law of partial pressures, that 
when two gases which do not act on each other are mixed, each 
exerts the same pressure as if it filled the whole space alone, and the 
total pressure is the sum of the pressures exerted by each gas. 


42 A TEXTBOOK OF CHEMISTRY 


Graphical Representation of the Gas Laws. It is often useful 
in studying physical and chemical phenomena to use a method 
of graphical representation which is illustrated in Figs. 10 and 
11. In Fig. 10 distances from the line OX represent pressures, 
while distances from the line OY represent volumes. If we 


VOLUMES 





PRESSURES 
Fig. 10 


start with a unit volume of a gas under unit pressure, represented 
by the point a, as the pressure increases the volume will decrease 
along the line ac, while as the pressure decreases the volume will 
increase along the lines ab and PV, the product of pressure and 
volume must always remain constant. The geometrical 
curve which satisfies these conditions is a hyperbola. 

Charles’s law may be represented in a similar way by Fig, 11. 


LAWS OF GASES 43 


Here the relation between volumes and absolute temperatures 
is represented by a straight line, but all gases liquefy before the 


VOLUMES 





ABSOLUTE TEMPERATURES 
Pig Le 


absolute zero is reached, and the line can never be continued, 
* experimentally, to the origin. 


EXERCISES 


1. A quantity of gas fills a volume of 175 ec. at 20° and under a 
pressure of 735 mm. What will be its volume under standard con- 
ditions (0° and 760 mm.) ? 

2. A flask having a capacity of 3.5 liters is filled with oxygen at 
25° and 770 mm. What weight of oxygen does it contain? (See 
p. 22.) 


3. What volume will 33.5 cc. of a gas measured at 18° and 715 mm. 
assume at 25° and 731 mm. ? 


4. A cylindrical gasometer has a diameter of 30 cm. and height of 
60 cm. What weight of oxygen will be required to fill it at 22° and 
745 mm. ? 


5. A bulb having a capacity of 127.2 cc. was exhausted till the 
manometer showed a pressure of 35 mm. while the temperature was 


1 The line OY is called the azis of abscissas and any line parallel to 
it and perpendicular to OX is called an abscissa, while OX is the azis 
of ordinates and any line parallel to this and perpendicular to OY 
is called an ordinate. O is called the origin. 


44 A TEXTBOOK OF CHEMISTRY 


23°. After weighing, it was filled at atmospheric pressure with a gas. 
The reading of the barometer was 751 mm. and the temperature 23°, 
as before. The increase in weight was 0.1382 gram. What is the weight 
under standard conditions of one liter of the gas which was used? 

(Ans. 1.2507.) 


6. A sample of bituminous coal has the following composition : 


Carbon 75.00 per cent 
Hydrogen 5.25 per cent 
Oxygen 10.00 per cent 


Ash, nitrogen, ete. 9.75 per cent 
100.00 per cent 


What is the heat of combustion of the coal in calories per kilo- 
gram and in B. T. U. per pound, assuming the heat of combustion 
of one gram of carbon as 8080 calories, one gram of hydrogen as 
34,179 calories (burned to liquid water), and that 1.25 per cent of the 
hydrogen is combined with the oxygen and contributes ‘nothing to 
the heat of combustion? It is assumed further that the combination 
between the carbon and hydrogen is of such a nature that these ele- 
ments give the same amount of heat when the coal is burned as they © 
would give if they were in the free state. 


CHAPTER IV 


HYDROGEN 
SymBoL, H.’ Atomic WeigGuHT, 1.0078. 


Occurrence. Although the quantity of hydrogen in the world 
is very much smaller than the quantity of oxygen, it is very 
widely diffused, especially in the form of its most common com- 
pound, water. It forms a little more than one ninth of the 
weight of water and is present both as water and as a constitu- 
ent of all of the most important compounds found in vegetables 
and animals. Hydrogen is an essential element, also, in the 
large class of substances called acids. A minute quantity, pos- 
sibly 0.001 per cent, or one part in 100,000, is found free in the 
air. ‘There is some evidence that at very high altitudes the at- 
mosphere consists almost exclusively of hydrogen. 

Acids. In order to understand one of the most convenient 
methods for the preparation of hydrogen in the laboratory, it 
is necessary to know something of the properties of the important 
class of substances called acids. It has been shown that when 
sulfur is burned in the air, sulfur dioxide, SOs, is formed. A 
small amount of the sulfur usually, or perhaps always, combines 
with more oxygen to form sulfur trioxide, SO3, and by means of 
suitable apparatus and a catalytic agent, nearly all of the sulfur 
can be converted into this compound (seep. 175). When sulfur 
trioxide is dissolved in water it combines with it, giving sulfuric 
acid : 

SO; + H,O = H,SO, 


Sulfur Water Sulfuric 
Trioxide Acid 


It will be recalled that this sort of combination between oxides 


of nonmetallic elements and water gave to chemistry the name 
45 


46 A TEXTBOOK OF CHEMISTRY 


of oxygen. Sulfuric acid, when pure, is a heavy liquid of 
an oily consistency, sometimes called oz! of vitriol. 

By the action of sulfuric acid on common salt we can obtain 
hydrochloric acid, HCl, a gas which dissolves easily in water and 
which is ordinarily used in the form of its solution. 

By the action of sulfuric acid on saltpeter, nitric acid, HNOs, 
is formed. This is a liquid, which is usually diluted with water 
for use. 

The most common acid of ordinary experience is acetic acid, 
HC2H;Os, the acid of vinegar. This is the acid from which we 
have all learned to associate the word acid with the sour taste 
which is characteristic of all moderately strong acids. 

Radicals. An examination of the formulas of the acids 
named above shows that each of them contains hydrogen, but a 
still more important characteristic of these and of all other acids 
is that in a great variety of reactions this hydrogen may be re- 
placed by other elements and especially by metals. The follow- 
ing are illustrations of such replacement : 


H.SO, + Zn == ZnSO, + 2 H 
Sulfuric Zine Zine Sulfate Hydrogen 
Acid 
HCl at Nar. NaCl + H 
Hydrochloric Sodium Sodium Chloride 
Acid (Common Salt) 
HNO; + NaOH = NaNO; + H,O 
Nitric Acid Sodium Sodium Water 
Hydroxide Nitrate 
HC.H;0. + NaOH = NaC.H;O0. + HOH 
Acetic Sodium Sodium Water 
Acid Hydroxide Acetate 


In each of these reactions one or two atoms of hydrogen are 
replaced by an atom! of some metal, while all of the rest of the 


1 One of the most common mistakes of beginners in such cases 
is to say ‘‘one or two parts of hydrogen are replaced by one part 
of the metal.’’ The distinction between one part and one atomic 
weight (or in accordance with the atomic theory, one atom) of an 
element ought always to be kept clear. 


HYDROGEN 47 


acid passes into the new compound without any change in com- 
position. A group of atoms, which remain in combination in 
this way when they pass from one compound to another, is 
called a radical. Thus SQ, is the radical of sulfuric acid ; NOs, 
of nitric acid ; C.H Os, of acetic acid. 

Salts. The compounds formed by the replacement of the 
hydrogen of an acid by a metal are called salts. These are so 
intimately connected with the acids in their composition that it 
is natural to use for them names which are derived from the 
names of the acids. How this is done is clear from the illustra- 
tions given. The name of the metal of the salt is placed first 
and this is followed by a word in which the -ic of the acid is 
changed to -ate. Sulfuric acid gives sulfates; nitric acid, nitrates ; 
acetic acid, acetates. The name of common salt, sodium 
chloride, seems to be an exception, but this 1s because, as a binary 
compound, it belongs to the class of substances which take 
names ending in -ide (p. 29). Additional principles which are 
used in naming acids and salts will be considered later. 

Preparation of Hydrogen. 1. Electrolysis of Dilute Sulfuric 
Acid. If an electrical current is passed between two strips of 
platinum (called electrodes) which are immersed in dilute sul- 
furic acid, bubbles of gas will rise from the electrodes ; and if an 
apparatus is so arranged (p. 9) that these can be collected, 
it will be found that the gas rising from the negative electrode 
(cathode) is hydrogen, while that from the positive electrode 
(anode) is oxygen. The volume of the hydrogen will be almost 
exactly twice that of the oxygen. , 

Since, as we shall find later, the hydrogen and oxygen are lib- 
erated in the same proportion in which they combine to form 
water, this experiment is often spoken of as a decomposition 
of water by electricity — and in a certain sense this is correct. 
That the sulfuric acid is more than a merely passive agent in 
what takes place is evident, however, first, because pure water 
is nearly a nonconductor for electricity ; and, second, because if 
we examine the liquid in the U-tube by appropriate means, we 
shall find that the hydrogen atoms of the sulfuric acid are being 


4S A TEXTBOOK OF CHEMISTRY 


transferred through the liquid toward the cathode as the current 
passes, while the radicals of the sulfuric acid, the SO, group of 
atoms, are transferred toward the anode. In other words, elec- 
trolysis is not merely something which takes place at the two 
electrodes, but it is always accompanied by a transfer of material 
through the whole of the space between ; and while hydrogen is 
carried in one direction, it is the sulfate radical and not oxygen, 
which is carried the other way. 

Electrolytes. Ions. Theory of electrolysis. Any substance 
which carries the electric current in this way is called an elec- 
trolyte. 'The most satisfactory theory which has been proposed 
‘to explain the facts which have just been given is that electro- 
lytes in solution are more or less completely separated into parts 
which are charged with positive or negative electricity. Ac- 
cording to this theory sulfuric acid separates partly into hydro- 
gen atoms with a positive charge of electricity and the sulfate 
radical with two negative charges. This is indicated by the 
symbols H+, H*, SO;-~. When the positive and negative elec- 
trodes are dipped in the dilute acid, the positively charged hy- 
drogen atoms are attracted by the negative cathode and move 
toward it, while the negatively charged sulfate radicals are re- 
pelled by the cathode and attracted by the positive anode. 
_ This causes the motion of the hydrogen atoms in one direction 
and the motion of the sulfate radicals in the other, through the 
solution. This motion constitutes the current of electricity in 
an electrolyte. At the cathode the hydrogen atoms lose their 
positive charge and at once appear as hydrogen gas. At the 
anode the action is more complicated, but the final result is 
that oxygen gas is liberated. The charged atoms or groups 
are called tons. The positive ion is called the cation, the neg- 
ative ion, the anion. ‘The decomposition of an electrolyte by 
an electric current is called electrolysis. 

2. Preparation of Hydrogen from Iron and Steam. If steam” 
is passed over red-hot iron contained in an iron tube (Fig. 12), 
a part of it gives up its oxygen to the iron and hydrogen is 
liberated. If the compound which remains in the tube is ex- 
amined, it is found to have the same composition as the magnetic 


HYDROGEN 49 


oxide of iron, Fe;04, formed when iron burns in oxygen. The 
equation is not quite so simple as those which have been given 





before. In order to arrive at the true equation the formulas of 
the substances used and the products obtained should be written 


first : 
Fe + H,O a Fe;O, — H 


Iron Water Magnetic Hydrogen 
Oxide of Iron 


On examining the above it is seen that 4 atoms of oxygen will 
be required to form one molecule of the magnetic oxide of iron, 
hence we must have 4 molecules of water in the first member of 
the equation to furnish these. The 4 molecules of water will 
give 8 atoms of hydrogen and 3 atoms of iron will also be re- 
quired to form the magnetic oxide. Putting all together we 
have: 


3 Fe + 4H.O = Fe;0, + 8H 


It would, doubtless, be easier for a beginner to learn this last 
equation outright than to learn how to, derive it in the manner 
indicated, but things which are a mere matter of memory are 
likely to be evanescent, while a rational process like the above 
can be reproduced at will. It is very important in studying 
chemistry to reduce those portions which are remembered as 
distinguished from those portions which are derived by a logical 


50 A TEXTBOOK OF CHEMISTRY 


process just as far as possible. At the same time many simple, 
fundamental facts, as here the composition of the magnetic 
oxide of iron, must be remembered and used over and over again. 

Reversible Reactions. It was stated above that a part only 
of the steam is decomposed by the iron. If we reverse the con- 
ditions and pass hydrogen over magnetic oxide of iron, part of 
the hydrogen will be converted into water and metallic iron will 
be obtained. The earlier and most natural idea of chemical 
affinity was that when three elements are present those two 
would unite which had the strongest affinity for each other. 
If this were true, either the hydrogen would be able to take the 
oxygen from the iron or the iron should be able to take it from the 
hydrogen. We see from the experiments described that this 
simple idea is not correct, but that either element can take the 
oxygen from the other. While there is a certain sense in which 
iron has a stronger affinity for oxygen than hydrogen has, the 
direction of the reaction depends on the quantities of the sub- 
stances present as well as upon their relative affinities. If 
steam is used and the hydrogen is constantly removed, the tend- 
ency is to form magnetic oxide of iron and hydrogen. If hydro- 
gen is used and the steam is constantly removed, the tendency 
is to form metallic iron and water. 

It is often convenient to express such reversible reactions as 


tollows 3 Fe + 4H,O => Fe,0, + 8H 


3. Decomposition of Water by Metals at Ordinary Tempera- 
tures. Potassium and sodium have a much stronger affinity 
for oxygen than iron has, and partly for this reason, partly, 
perhaps, for other reasons which we do not fully understand, 
these metals will decompose water and liberate hydrogen at 
ordinary temperatures. . If potassium is thrown on water, the 
heat of the reaction is great enough to cause the hydrogen to 
ignite. It burns with a violet flame, the color being given to it 
by the potassium. Sodium when thrown on water usually rolls 
over the surface in a globule, evolving hydrogen, which does not 
take fire, but if thrown on a piece of filter paper lying on the 


HYDROGEN fal 


water so that the globule remains at one spot, the hydrogen will 
catch fireand burn with the yellow flame characteristic of sodium. 
If a small piece of sodium is wrapped in paper and thrust quickly 
under the mouth of a jar which has been 
filled with water and inverted with the 
mouth under water, the sodium will act on 
the water as before and the hydrogen may 
be collected and examined. 

If the water in which the potassium or 
sodium has dissolved in these experiments is 
examined, it will be found to have a soapy 
feel and disagreeable, acrid taste. It will 
also turn the color of red litmus paper blue. 
If the solution is evaporated in a dish of Fig. 13 
platinum or of some material which is not 
affected by it and under such conditions that it cannot absorb 
carbon dioxide from the air, a white solid will be obtained, 
which will have the composition represented by the formula 
KOH or NaOH. These substances absorb and retain water 
so greedily that it is necessary to heat them nearly to redness 
before the last of the water is expelled. They are called, in 
accordance with their composition, potassium hydroxide or 
sodium hydroxide. The equations are: 





K + HO = KOH aa H 
Potassium Water Potassium Hydrogen 
Hydroxide 
Na + HO = NaOH + H 
Sodium Sodium 
Hydroxide 
or, Na + HOH = NaOH + H 
Hydrogen 
Hydroxide 


Contrast between the Action of Iron and of Sodium on Water. 
The last form expresses a little more clearly that the metal has 
replaced only one of the two atoms of hydrogen in each molecule 


52 A TEXTBOOK OF CHEMISTRY 


of water. The action is seen to be quite different from that of 
iron on steam. ‘This is closely connected with the amount of 
chemical energy changed to heat in each reaction. In the reac- 
tion, Na + H,O = NaOH + H, 43,450 calories are liberated 
for each gram atom of hydrogen set free, while in the reaction, 
3 Fe + 4 H.O = Fe30, + 8 H, if it could be carried out at 100°, 
only 4160 calories would be given for each gram atom of hydro- 
gen liberated. In general, those reactions in which large 
amounts of chemical energy are changed to heat take place most 
easily. We must, however, guard against the impression that 
this is a universal law. The ease with which a reaction takes 
place is by no means proportional to the heat generated. Other 
factors are involved, and some of these are, at present, but little 
understood. 

4. Hydrogen from “ Hydrone.”’ The action of water on so- 
dium is too violent for use as a laboratory method of preparing 
hydrogen in quantity. If, however, the sodium is alloyed with 
lead, the action is moderated, and such an alloy containing about 
35 per cent of sodium is sold under the name of “ hydrone.”’ 
By means of it very pure hydrogen can be easily prepared. 

}. Preparation of Hydrogen by the Action of Metals on Acids. 
If a strip of zine and one of platinum, copper or lead are dipped 
in dilute sulfuric or hydrochloric acid while the strips are con- 
nected by means of a wire, an electrical current will pass through 
the wire while bubbles of hydrogen will be seen to collect and rise 
from the surface of the platinum, copper or lead. If the liquid 
between the two metallic plates is examined, as in the electrolysis 
of dilute sulfuric acid, it will be found that the hydrogen travels 
through the liquid toward the platinum, while the sulfate radical 
or the chlorine travels toward the zine. -At the surface of the 
zinc, the sulfate radical or the chlorine combines with the zine, 
forming zine sulfate, ZnSO, or zine chloride, ZnCle. If pieces 
of chemically pure zine are placed in dilute hydrochloric or sul- 
furie acid, there will be almost no action at all, while commercial 
zine will dissolve rapidly. After the action of the acid on the 
commercial zinc has continued for a short time it will be seen 


HYDROGEN 53 


that the surface is dark, and a closer examination will show that 
it is covered with lead and other impurities found in the zinc. 
When we consider these facts along with the experiment with the 
strips of platinum and zinc, we reach the conclusion that the ac- 
tion of the acid on the zine requires some catalytic agent like 
the platinum or lead before it can be very rapid, and that the 
phenomenon of the solution of the zinc is partly electrical, being 
accompanied by electrical currents between the particles of lead 
and zinc in the impure zine. If we disregard the catalytic agent, 
the process may be expressed by the equations : 


Zn + H2SO, = ANDO, ee ok 
Zine Sulfuric Zine Hydrogen 
Acid Sulfate 
Zn + 2 HCl =i /nUlse aye ° 2) H 
Hydrochloric Zine 
Acid Chloride 


It is to be noticed that one atom of zinc replaces two atoms of 
hydrogen in each case and that when an acid is used which has 
only one atom of hydrogen in its molecule, two molecules of the 
acid are required for the reaction. 

If iron is substituted for zinc, these reactions become: 


Fe ot H.SO, => FeSO, f 2 H 
Tron Sulfuric Ferrous Hydrogen 
Acid Sulfate 
Fe + 2HClL = FeCl + 2H 
Ferrous 
Chloride 


Apparatus for the Preparation of Hydrogen. In the labora- 
tory, small quantities of hydrogen may be generated in the 
simple apparatus shown in Fig. 14. Zine and some water are 
placed in the generating flask, and dilute sulfuric acid! is added 
in portions through the thistle tube. 

1 As sulfuric acid is heavier than water (sp. gr. 1.84) and much 
heat is generated on its dilution, it should always be poured slowly 
into water and should never be diluted by pouring water upon the 


acid. Pouring water on concentrated sulfuric acid may cause an 
explosion. 


o4 


A TEXTBOOK OF CHEMISTRY 


A more convenient apparatus for the preparation of larger 
amounts, or when it is desired to have the gas always ready for 





0 By 
vee 7-8 Sag 
Frese 


235 tt -— 


Fig. 14 





use, is the Kipp generator (Fig. 15). The 
zinc is placed in the middle bulb and the 
dilute acid is poured in through’ the upper 
bulb, which communicates with the lower one 
through the tube A. When the stopcock B 
is opened, the acid rises and comes in contact 
with the zine in the middle bulb and the gen- 
eration of hydrogen begins. Whenever the 
stopcock is closed the hydrogen generated 
forces the acid away from the zine and the 
action ceases as soon as the acid moistening 
the surface of the zinc is exhausted. The 


generator is not altogether satisfactory because the spent acid 
containing zine sulfate is mixed with that which has not been 


used, diluting it and causing the 
action to become very slow before 
the acid has been exhausted. A 
more suitable form of apparatus 
for generating large quantities of 
hydrogen is described on p. 165. 
Purification of Hydrogen. The 
hydrogen prepared by any of the 
methods. described is impure. 
Spray from the generating liquids 
may be removed by passing the 
gas through a tube filled with cot- 
ton wool. Moisture, or water 
vapor may be removed by means 
of calcium chloride,! contained in a 


1Qne liter of a gas dried with cal- 
cium chloride_retains 0.36 mg. of water 
at 25°. When dried with concentrated 





Fig. 15 


sulfuric acid the amount of water retained by one liter of the gas is only 
0.002 mg. at 15° to 19°. When dried with phosphorus pentoxide 


HYDROGEN 


55 


tube such as shown in Fig. 16, or more perfectly by means of 
pumice stone or glass beads moistened with concentrated sul- 


furic acid or by phosphorus pentoxide. Hydrogen 
sulfide ! and some other impurities, especially some 
of those wpich give an unpleasant odor to the gas, 
may be removed by passing it through a wash bottle 
containing a solution of potassium permanganate, 
but a small amount of oxygen will be introduced 
into the gas (V. Meyer, and Recklinghausen Ber. 29, 
2550). Oxygen may be removed by passing the gas 
through a hot tube containing platinized quartz, 
which will cause the oxygen to combine with some of 
This should be done, of course, be- 
fore the gas is dried. Nitrogen 
from the air cannot be removed 
and when pure hydrogen is re- 
quired very great care is necessary 
to prevent its entrance (Cooke 


and Richards, Am. Chem. J. 10, 


the hydrogen. 








y 


ye 


Chem. Soc. 30, 1724). 
Properties of Hydrogen. 


\ 
to) 
' 








colorless, tasteless and odorless gas. 









Yass aro 


‘> 
t 


oes): 
2. 
@, 






eos 

S 2 Cahn gy! a 
eae teee 

ah SOTA SO) 


a 
2 
her 


wey, 
ey 
1 2, 






& LO a 
ie 38a 
Mattos 





oe 
D(a GT 
Yast @ 





PRX 


ex 


81; Morley, chad. 17, 267; Noyes, J. Am. 


Hydrogen is a 


One 





280% ‘ U 
722000000 99900 





liter weighs at 0° and 760 mm. pressure 
0.08987 gram. As the weight of a liter of 





—s Pia leva 
1.293 : Ree VY 
air is grams, air is 7 icog7 
Fig. 17 times heavier than hydrogen, or approxi- 
mately 145 times. Oxygen is is 2 15.90 times heavier 
0.08987 


than hydrogen, or approximately 16 times. 


the amount of moisture retained by one liter of the gas is less than 


0.00002 mg. (Morley). 


1 Hydrogen sulfide may be removed to better advantage by 
passing the gas through a wash bottle or serpentine tube containing 


lead oxide dissolved in a solution of potassium hydroxide. 


56 A TEXTBOOK OF CHEMISTRY 


At a very low temperature hydrogen condenses to a liquid 
which boils under atmospheric pressure at — 252.5° or 20.5°, ab- 
solute. The liquid has a density of only 0.07 gram per cubic 
centimeter or one gram fills a volume of about 14 cc. If the 


Fig. 18 


liquid is made to boil by reduc- 
ing the pressure, it grows still 
colder and at — 260°, or 13°, ab- ° 
solute, what remains freezes to 
a solid. The vapor pressure of the 
solid is 58 mm. 

The lightness of hydrogen as 
compared with air may be easily 
shown by pouring it upward 
through the air, by showing that 
it may be collected in an inverted 


jar or beaker while it will not remain in one which is placed ° 
upright, and by filling soap bubbles or toy balloons with the gas. 


The use of the gas for filling balloons is well known. 
* What is the lifting power of one cubic meter 
of pure hydrogen under standard conditions? 
What will be the lifting power of one cubic meter 
at a mile above sea level when the pressure 
is 620 mm. and the temperature 20°? It is con- 
venient to remember that as the volume of a gas 
varies inversely with the pressure and directly 
with the absolute temperature, the weight of a 
given volume must vary directly as the pressure 
and inversely as the absolute temperature. 
Diffusion of Gases. _If two cylinders are filled, 
one with hydrogen and the other with air, and 
placed with their mouths together but with the 
hydrogen above (Fig. 19), it will be found after a 
comparatively short time that the gas in each 


cylinder will explode, if ignited, with the whistling sound char- 
acteristic of a mixture of air and hydrogen fired in an open 
cylinder or test tube. Although the air is fourteen and a half 


Fig. 19 


DIFFUSION 57 


times heavier than the hydrogen, it makes its way quite rapidly 
up into the hydrogen above it and the hydrogen passes down 
into the heavier air below. This property of mixing with each 
other is true of all gases without exception and is called 
diffusion. While there are many liquids which do not dissolve 
in each other or which dissolve only to a limited extent and 
such liquids separate into layers in accordance with their 
specific gravities, gases, which differ much more than liquids in 
their densities, always mix when brought 
in contact and when once mixed will sep- 
arate to only a very slight extent (p. 45) 
in accordance with their densities. Very 
closely related to the diffusion of gases 
which are in contact with each other is 
the diffusion of gases through a wall full 
of fine openings, which separates them. 
If a cylinder of porous porcelain with 
openings so fine that pressure will cause a 
gas to pass through them only very slowly 
is fitted with a rubber stopper and con- 
nected with a bulb and bent tube filled 
with water as shown in Fig. 20, on bring- 
ing a beaker filled with hydrogen over 
the cylinder the pressure within will sud- 
denly increase and force water out of the 
tube in a jet, showing that hydrogen is 
passing through the walls of the cylinder to the interior. 
When the stream of water ceases, if the beaker is removed, 
the movement of the water in the tube will show that diffu- 
sion of the hydrogen outward is taking place. By appro- 
priate experiments it is possible to show that some air passes 
out through the porous wall while the hydrogen is passing in. 
A careful study of the phenomena by Graham has shown that 
gases pass through a porous wall of this sort, at a rate which 
varies inversely as the square root of the density. Oxygen, 
which is about sixteen times as heavy as hydrogen, will pass 
the wall only one fourth as fast. 


coe 
jy 
ZZ Z 


Wy 

wy 
GEA 

——— 





AN 
Wy 


=(])——_ p= 
N 
Re ) 


Fig. 20 


58 A TEXTBOOK OF CHEMISTRY 


Kinetic Theory of Gases. When water is converted into 
steam at 100° the volume at atmospheric pressure is increased 
more than 1600 times. This fact and many of the other proper- 
ties of gases makes it seem highly probable that the space be- 
tween the molecules of a gas is very large in comparison with the 
size of the molecules themselves. A very satisfactory explana- 
tion of the law of diffusion of gases, given in the last paragraph, 
and of the fact that a gas expands at once to fill any empty space, 
however large, which it is allowed to enter, is found in the kinetic 
theory of gases. According to this theory the molecules of 
gases are moving constantly at a comparatively high velocity, 
and whenever they meet each other they rebound according to 
the laws of elastic bodies. According to these laws when two 
elastic bodies, of equal weight, meet each other squarely each 
rebounds with the energy of the other. If the bodies meet at 
an angle, the interchange of energies will be only partial, and 
the effect of this constant interchange must be to give to all 
molecules of the same weight approximately the same average 
energy and hence the same average velocity. If molecules of 
different weights are mixed, however, the interchange of energies 
must give a greater velocity to lighter molecules. As the energy 
of a moving body varies as the square of its velocity, the veloci- 
ties of molecules of different weights must vary inversely as 
the square roots of their weights, if their energies are the same, 
since the pressure of a gas is due to the impacts of its molecules 
on the walls of the containing vessel and the pressure does not 
change when two gases having different densities are mixed. 

As an illustration we may take oxygen and hydrogen. The 
molecule of oxygen is 16 times as heavy as that of hydrogen 
(p. 95). If the two gases are mixed, the molecules of hydrogen 
must, through frequent collisions with molecules of oxygen, soon 
have the same average energy as the latter, and to do this must 
have, on the average, four times the velocity of the molecules 
of oxygen. What is true of the mixed gases must be true also 
of the gases when separate. Accordingly, if we have a porous 
wall with fine openings separating the two gases, as the hydro- 


DISSOCIATION 59 


gen molecules have four times the velocity of the oxygen mole- 
cules and there are the same number in equal volumes, four 
hydrogen molecules will hit the openings and pass into them 
while one oxygen molecule does so. This ratio of four to one 
is the same as the ratio of the square roots of the densities, 
V16: V1 = 4:1, the law for diffusion through a porous wall 
given above. 

Chemical Properties of Hydrogen. If hydrogen is brought to 
the air, through a glass tube drawn to a narrow opening, and 
lighted, it will burn for an instant with a pale blue, almost in- 
visible flame, but the color quickly changes to yellow from par- 
ticles of sodium or its compounds, which are volatilized from the 
glass by the heat. From a platinum jet the pure gas burns with 
a flame almost or quite invisible in daylight. The product 
formed is water, as may be shown, roughly, by holding a cold 
glass over the flame, or, more accurately, by burning the gas for 
some time, condensing the water formed and determining its 
freezing point and boiling point. The hydrogen and air used 
should, of course, be carefully dried. 

Mixtures of air and hydrogen explode when ignited, as the 
flame travels through such a mixture with a very high velocity, 
and both the steam formed and the nitrogen 
of the air are heated to a high temperature 
and expanded greatly by the heat of com- 
bustion. Mixtures of oxygen and hydrogen 
explode still more violently. Hydrogen does 
not support the combustion of substances 
which burn in oxygen (Fig. 21). It will 
support the combustion of oxygen or chlo- 
rine. 

Dissociation. If steam is passed through 
a tube of porous porcelain A, Fig. 22, which 
is inclosed in a larger tube B of glazed porce- 
lain through which is passed a current of 
carbon dioxide, introduced through C, while the whole is 
heated to a very high temperature (2000°, perhaps) the steam 





Fig. 21 


60 A TEXTBOOK OF CHEMISTRY 


will be partly decomposed into oxygen and hydrogen, and, 
since the hydrogen diffuses through the porous porcelain four 
times as fast as the oxygen, more oxygen than enough to com- 
bine with the hydrogen which remains will stay in the inner 
tube while more hydrogen will pass through than enough to 
combine with the oxygen which passes through. If the cooled 










Io 
v 


8 
. 
G io 






os 
AS 
& 





soho Ly 9 
“oO 





gases which are delivered at the ends of the tube are mixed 
and the carbon dioxide is absorbed by a solution of sodium hy- 
droxide, it will be found that the gas which remains consists 
of a mixture of two volumes of hydrogen with one volume of 
oxygen. In this way Deville showed that water can be de- 
composed into oxygen and hydrogen by heat alone, and that 
the reaction between hydrogen and oxygen is reversible : 


2H+O0SH,0 


A decomposition of this kind, when the products of decomposi- 
tion recombine on cooling or on a reversal of the process which 
caused the decomposition, is called a dissociation. Probably all 
compounds would be decomposed into their elements at a suffi- 
ciently high temperature, but it is only in those cases where the 
elements recombine, on cooling, to form the same compound, 
that the decomposition is called a dissociation. 

* By other methods the per cent of dissociation of water has 
been determined up to 2300°, absolute. From the results the 
dissociation at still higher temperatures may be calculated 


OXYHYDROGEN BLOWPIPE 61 


approximately.|. The results are as follows, at atmospheric 
pressure : 














ABSOLUTE TEMPERATURE ee ts CENTI- Reha OF 
ie” i 
1000° (Pane 0.00003 
1500° 122 fe 0.022 
2000° Lita 0.59 
2500° hy Me 3.98 
2113. 2500° 8.12 
B2ipe 3000° 20.0 
4000° eects 40.5 











The Oxyhydrogen Blowpipe. If oxygen and hydrogen are 
brought together in sucha manner that they burn as they come 
in contact, an almost colorless flame having a very high tempera- 





ture is produced. If steam could be heated to very high tem- 
peratures by the expenditure of the same quantity of energy, pro- 


1 By the formula, 
Boke (LOU ee 11.46 _ 25,030 2 38 ] We 
CER eeciteey Cian ean ae TR ee LOGO 
— 1.38.104(7 — 1000) — 0.685.10-7(72 — 10002), in which P is 
the pressure in atmospheres, 7, the absolute temperature, and z, 
pe. Sh a dissociated. (Nernst, Theoretische Chemie, 6% Aufl., 
p. é 








62 A TEXTBOOK OF CHEMISTRY 


portionally, as that required to heat it to 1000°, the heat of com- 
bustion of oxygen and hydrogen (p. 65) is great enough to heat 
the steam produced to 10,000°, at least. A little consideration 
of the dissociation of water at high temperatures shows us that 
such an extreme temperature cannot be reached, for it is evident 
that oxygen and hydrogen cannot combine to produce heat at 
a temperature at which water largely decomposes into its ele- 
ments. Since there is a dissociation of 40 per cent at 3700°, it is 
doubtful if even that temperature can be obtained. ‘The tem- 
perature of an electric arc between carbon points is estimated 
at about 3600°. The temperature of an open-hearth steel fur- 
nace is only 1500° to 1700°. 

Iron wire will take fire in the oxyhydrogen flame and burns 
brilliantly. Platinum melts easily (1755°) and the flame has long 
been used in working with this metal. A piece of lime held in 
the flame glows intensely, giving the light known variously as 
the oxyhydrogen, lime or Drummond light. It will be noticed 
that any solid substance, which does not volatilize too easily, 
gives an intense light in the flame, while the oxyhydrogen flame 
alone is almost nonluminous. 

Explosions. Catalysis. A mixture of oxygen and hydrogen 
may remain in a glass tube indefinitely without combining to an 
extent that can be measured. It is not till a temperature of 
300° is reached that the gases combine rapidly enough so that 
the rate of combination can be measured after some days or 
weeks. At a temperature a little above 500°, the combination 
is fast enough so that the heat of combination raises the mix- 
ture to a higher temperature, which still further accelerates the 
combination, and an explosion results. It is characteristic of 
most explosions caused by chemical action that the reaction 
causing the explosion is accelerated enormously by the heat 
of the reaction. If the mixture of oxygen and hydrogen is 
brought into contact with platinum in the spongy form or with 
platinized asbestos ! the reaction is hastened and the gases will 


1 Prepared by moistening asbestos with a strong solution of 
chloroplatinie acid and heating it. 


OXIDATION. VALENCE 63 


usually take fire and burn. This is another illustration of ca- 
talysis and recalls the effect of other metals on the solution of 
zinc in acids. It may be that the two phenomena are closely 
related. 

Oxidation. Reduction. If a piece of copper is held over a 
flame so that it is heated quite hot while exposed to the air, it 
will be oxidized, the surface changing to black copper oxide. 
On holding the hot, oxidized copper in an atmosphere of hydrogen, 
the black oxide will be quickly changed back to metallic copper. 
This process is called reduction, and hydrogen is called a reducing 
agent. It will be seen from the above that oxidation and reduc- 
tion are opposite processes. The two words are often used in 
chemistry in a much more general sense. The addition of other 
elements than oxygen is frequently called oxidation, and the re- 
moval of other elements, or the substitution of hydrogen for 
other elements, or even the addition of hydrogen to a com- 
pound, may be called a reduction. 

Valence. It may have been noticed that when sodium and 
zine act on water or on hydrochloric acid, one atom of the sodium 
replaces one atom of hydrogen, while one atom of zinc replaces 
two atoms. This characteristic of metals may be made more 
striking by selecting three metals whose atomic weights are close 
together. If 23 milligrams of sodium, 24 milligrams of magne- 
sium and 27 milligrams of aluminium are allowed to act on hy- 
drochloric acid in such a way that the hydrogen generated is 
collected in separate tubes,! it will be found that the sodium will 
give about 11 cc. of hydrogen,? the magnesium 22 cc. and the 
aluminium 33 ce. The property of the metals illustrated here 
is called valence. A metal which replaces one atom of hydrogen 


1 The sodium may be placed in a short lead tube, 38 mm. in diam- 
eter and closed at one end. The mouth should be closed with a 
little cotton till ready for use. The magnesium and aluminium may 
be weighed in small watch glasses about the size of the mouths of 
the test tubes to be placed over them. The aluminium must be 
etched with a solution of sodium or potassium hydroxide and after- 
wards washed and dried before use. 

2One ce. of hydrogen weighs 0.09 mg. 0.09 11.1= 1 mg. 
of hydrogen from 23 mg. of sodium. 


64 A TEXTBOOK OF CHEMISTRY 


for one atom of the metal is called wnivalent; one which re- 
places two atoms is called bivalent; three, trivalent; four, 
quadrivalent. On a somewhat different basis, which will be 
discussed later (p. 156), when other meanings of valence 
are considered, some elements may be quinquivalent, sexivalent, 
septivalent or even octovalent. The valence of the metals is 





Fig. 24 


such an important characteristic and a knowledge of it is so 
useful in the writing of formulas that it seems best to present 
it from the standpoint of replacement here. 

Some metals have two or more kinds of valence in different 
compounds. ‘Thus, in the sense in which valence is used here, 
iron is bivalent in ferrous chloride, FeCle, and ferrous sulfate, 
FeSO,, but trivalent in ferric chloride, FeCl3, and ferric sulfate, 
Fe2(SO.)3. This last formula illustrates the value of a knowledge 
of valence in writing formulas. (What is the formula of alu- 
minium sulfate? What are the formulas of stannous sulfate 
and of stannic sulfate if tin is bivalent in stannous compounds 
and quadrivalent in stannic compounds ?”) 

From the point of view of the atomic theory the facts by 
means of which the valences of elements are determined point 
very strongly to differences in the powers of atoms to combine 
with other atoms. Thus in the compounds NaCl, MgCl. and 


HYDROGEN 65 


AICI; it seems that an atom of sodium can hold one atom of chlo- 
rine in combination, an atom of magnesium can hold two and an 
atom of aluminium, three. This property of valence is often 
expressed by means of such formulas as the following : 


cl cl 
Na—Cl, Mg, AICI, jelag us 
cl ‘CI 


The lines are intended to represent lines of force holding the 
atoms in combination, and the number of lines proceeding from 
the symbol of an element indicates its valence. 

We must distinguish sharply between the intensity of the force 
holding two atoms together and the valence of the atoms. Thus 
the valence of sodium in Na—Cl is the same as that of hydrogen 
in H—Cl, but the force which holds the sodium and. chlorine 
together is much greater than that which holds hydrogen and 
chlorine together. 

Heat of Combustion of Hydrogen. ‘The heat generated when 
2.015 grams of hydrogen combine with 16 grams of oxygen and 
the water formed is condensed to the liquid state at 18° is 
68,414 small calories, or 34,179 calories for one gram of hydro- 
gen (Thomsen). This value is often used in calculating the 
heat of combustion of coal. 

If 2.015 grams of hydrogen combine with 16 grams of oxygen 
at 100° and the water formed remains as steam, the heat of 
combination is only 58,000 calories, at constant pressure, or 
28,970 calories for one gram. ‘This is, of course, the maximum 
amount of heat which can be obtained by burning hydrogen un- 
der practical conditions. 


CHAT DERRY, 


WATER, HYDROGEN PEROXIDE 


Analysis. Synthesis. The two methods by which the com- 
position of a substance is determined are by analysis, the sepa- 
ration of the substance into the elements of which it is com- 
posed, and by synthesis, the putting together of the elements to 
form the compound. In analysis it is comparatively seldom that 
the elements are separated in the free state. Thus, in order to 
determine the amount of hydrogen in an organic compound, 
such as sugar, the substance is burned and the water formed is 
collected and weighed. Knowing what per cent of hydrogen 
is contained in water, it is easy to calculate the amount of hydro- 
gen contained in the compound. Either an analysis or a synthe- 
sis may be qualitative, giving simply the elements which are 
present, or quantitative, giving the quantity or per cent of each 
element. 

Qualitative Analysis and Synthesis of Water. ‘The prepara- 
tion of hydrogen by passing steam over red-hot iron is a qualita- 
tive analysis of water. In order to make the analysis complete 
it would be necessary to show that hydrogen and the magnetic 
oxide of iron are the only products of the action, that the density 
and properties of the hydrogen are the same as those of hydrogen 
prepared in other ways and that the density and properties of 
the magnetic oxide formed are the same as those of the oxide 
formed by burning iron in oxygen. 

The experiment showing that water is formed when dry hy- 
drogen burns in air or in oxygen is a qualitative synthesis of 
water. 

Quantitative Synthesis of Water by Volume. The determina- 
tion of the composition of water by volume may be made in an 


instrument called a-eudiometer, a tube graduated usually to 
66 


COMPOSITION OF WATER 67 


tenths of a cubic centimeter by means of fine lines etched on 
the surface. For the synthesis of water two platinum wires 
must be sealed in, near the closed end. <A capacity of 50 ce. is 
suitable for the experiment to be described. Such a eudiometer 
is carefully dried and filled with mercury and 10-12 ce. of dry 
oxygen introduced. ‘The volume of the gas is then accurately 
measured and the height of the mercury in the eudi- 
ometer above the mercury in the reservoir, the tem- 
perature and the reading of the barometer are noted. 
From these measurements the volume of the oxygen 
under standard conditions is calculated. Enough 
dry hydrogen to give a total volume of 25-28 ce. is 
then introduced and these measurements repeated. 
The hydrogen should be in excess, but, as the pres- 
sure is greater, the total volume 
need not be three times the vol- 
ume of the oxygen. The eudiom- 
eter is then clamped firmly, with 
a piece of sheet rubber placed 
under its mouth, and the mix- 
ture exploded by passing an elec- 
tric spark between the platinum 
points by means of an induction 
coil. After cooling, the volume 


THT 


| 
D 


nu 


Lu 


—_ 
rs 


Steam. 





I} 
NR 
oOo 


I | 


JALAL 


(st) 
oS 


of hydrogen remaining is meas- 
ured as before, except that the 
water formed by the explosion 
remains partly as vapor in the 
hydrogen and the pressure of the hydrogen is 
to be found by subtracting from the reading 
of the barometer the height of the mercury 





Rigs 25 


in the eudiometer plus the pressure of vapor 

of water for the temperature which is read 

(see p. 75). After correction of the three volumes of gas to 

standard conditions the proportion by volume in which the gases 
ve united may be calculated. 


Fig. 26 


68 A TEXTBOOK OF CHEMISTRY 


By placing a tube over the eudiometer (Fig. 26) and passing 
steam through it, the water formed in the experiment may be 
converted into steam and the volume of the excess of hydrogen 
plus the steam determined and from this the volume of the steam 
calculated. : 

*The results of Morley’s exceedingly careful experiments 
(Amer. J. Sci. 41, 220 and 276; Chem. News, 63, 218) show that 
when oxygen and hydrogen are measured in tubes the ratio of 
the volumes which combine is, O: H = 1: 2.0002. Curiously 
enough Scott (Phil. Trans. 184, A, 543 (1893)) has found that 
when the gases are measured in globes the ratio is O:H = 
1: 2.0025. It seems, therefore, that the same quantity of gas 
may fill a different volume when measured in a tube from what 
it does when measured in a globe, but no one has proved this by 
direct experiment. 

The volume of the steam is very nearly the same as the volume 
of the hydrogen which goes to form it. We can express the re- 
lation by the following diagram : 


fe] + fal fe) = [ole 


1 vol. oxygen 2 vols. hydrogen 2 vols. steam 


Composition of Water by Weight. From the composition of 
water by volume and the weights of one liter of each gas we may 
calculate the composition by weight. This gives: 


O: H = 1.429: 2.0025 * 0.08987 
or 16:2 * 1.0075 


The Unit for Atomic Weights. The quantity of hydrogen 
combining with 16 parts of oxygen is given because an atomic 
weight of 16 has been assigned to oxygen, somewhat arbitrarily, 
as a basis for comparison with all other atomic weights. Hy- 
drogen with an atomic weight of one was originally chosen as 
the unit for atomic weights. For 70 or 80 years it was sup- 
posed, on the basis of inaccurate determinations of the composi- 
tion of water, that the atomic weight of oxygen was, on that 


COMPOSITION OF WATER 69 


basis, almost exactly 16 (or 8). When the composition of water 
was finally determined more accurately, it was decided by the 
majority of chemists that it is better to make oxygen, with an 
atomic weight of 16, the basis for all other atomic weights, rather 
than to make such large changes as would be necessary in many 
of the common atomic weights, if hydrogen were retained as 
the unit. A few chemists, however, still prefer hydrogen as the 
unit. , 

Determination of the Composition of Water by the Use of 
Copper Oxide. ‘The first moderately accurate determination of 
the composition of water was made by the Swedish chemist, 
Berzelius, in 1819. He weighed a quantity of copper oxide in 
a glass bulb, heated it, passed dry hydrogen through the bulb, 
and collected and weighed the water formed. The copper oxide 
was reduced to metallic copper, and the loss of weight gave the 
weight of oxygen which had been converted into water. The 
difference between the weight of the water collected and the 
weight of oxygen taken from the copper oxide gave the weight 
of the hydrogen. A number of years later (1842) a French 
chemist, Dumas, carried out an elaborate series of experiments 
by the same method, with the.apparatus shown in Fig. 27. 
The hydrogen was generated in the large bottle and passed 
through a series of tubes to purify and dry it. It was then 
passed through the bulb containing the copper oxide and the 
water formed was collected, partly in a bulb, and partly in dry- 
ing tubes. The atomic weight of hydrogen calculated from the 
results of 19 experiments with this apparatus is 1.0025.! For 
some unknown reason the result is too low by about one part 
in 200. About 50 years later the copper oxide method was 
modified by the author, who used the apparatus shown in Fig. 
28. After placing some copper oxide in the bulb A and ex- 


1The student is not, of course, expected to remember these 
various values. A very brief description of 4 out of some 16 deter- 
minations of the composition of water is given as an illustration of the 
amount of labor which has been expended on the determination of 
atomic weights and also to illustrate how successive workers attain, 
sometimes, a closer approximation to the truth. 


70 


A TEXTBOOK OF CHEMISTRY 























dill 


Raa 








































































































: 
i 


; 
_sl/ 
{ 
BE fe : 











SSS 
it <= 
= 
oP ei 
ws pre) 
. SS > 
Ss @& bh wb } Hoos ob ee SWZ oa 
Fa Ss Be & &§ Ps «OBS me FR ! ae 
= ci) fod be) 
i] i i) i g) 
f | f hig [ 













































































































































































































































































































































































































































































































































































































































































Fig. 27 


COMPOSITION OF WATER 71 


hausting it with a good mercury air pump the apparatus was 
weighed. It was then connected at C with an apparatus fur- 
nishing pure hydrogen, the bulb A was heated in an air bath and 
the tube B was cooled. ‘The stopcock E permitted the hydrogen 
to pass out at first through D ey no air should enter the bulb 
from the connecting tubes. On admitting hydrogen to the bulb 
it was converted into water by the copper oxide and the water 
was condensed in B. After fromjone to two grams of hydro- 





Fig. 28 


gen had been converted into water in this way the stopcock was 
closed and the apparatus cooled and weighed. The gain in 
weight was the weight of the hydrogen which had entered. The 
apparatus was then connected with a tube into which the water 
-formed could be driven by warming B and A. The loss in weight 
of the apparatus gave the weight of the oxygen which had been 
taken from the coppcr oxide. The water was also collected and 
weighed. Twenty-four determinations, partly by this method, 
partly by another which need not be described here, gave 
1.00787 as the atomic weight of hydrogen. 

Determination of the Composition of Water by weighing 
Oxygen and Hydrogen. For more than twelve years Professor 
Morley worked at Cleveland on the composition of water by 
volume, on the determination of the weights of oxygen and hy- 
drogen gases and finally on the composition of water by weight. 
The weights of one liter of hydrogen and of oxygen which have 
been given are the values determined in this long, classical in- 


72 . A TEXTBOOK OF CHEMISTRY 


vestigation. The composition of water by weight was also de- 
termined in the apparatus shown in Fig. 29. Into this apparatus 
were brought hydrogen from a tube containing metallic palla- 
dium, in which it had been absorbed, and oxygen from globes in 
which it had been weighed. ‘The water formed was frozen in 
the bottom of the apparatus and was weighed 
at the end of the experiment. From twelve 
experiments the atomic weight of hydrogen is 
calculated as 1.00762. The atomic weight of 
hydrogen which is now used (J. Am. Chem. 
Soe. 31, 1) is 1.0078. This is probably not in 
error by so much as one part in 5000. For 
ordinary calculations the value is rounded off 
to 1.008, or, frequently, to 1.01. 

Properties of Water. Pure water appears, 
ordinarily, to be colorless and transparent, but 
light transmitted through a layer of water some 
meters in thickness has a blue color. Water 
is the only substance for which we have three 
names according as it is in the form of a solid, 
liquid or gas. Water has a maximum density 
at 4°. If either cooled or heated from that 
temperature, it expands. For this reason in- 

Fig. 29 the fall and winter large bodies of water cool 

by convection, that is by the sinking of the 

cooler, heavier water on the surface, till a temperature of 4° is 

reached. On further cooling the water grows lighter again and 

the colder, lighter water floats on the surface, protecting the 

warmer water beneath from further rapid cooling. ‘The ice which 

finally forms has a density of only 0.92 (accurately 0.91674) and 

continues to float on the surface. The density of water at 
different temperatures is given in the following table: 1 



































1 This table is useful especially for determining the capacity of 
burettes or flasks by weighing the water which they contain. Thus 
it is seen that the water which will fill 1 ec. at 20° weighs 0.99823 
gram. ‘This is, however, when weighed in a vacuum. If weighed 
with brass weights in air, the apparent weight will be 0.00105 gram 


PROPERTIES OF WATER 








TEMPERATURE 


CO CONT OD Or PwNwe © 





DENSITY 


0.99987 
0.99993 
0.99997 
0.99999 
1.00000 


0.99999 
0.99997 
0.99993 
0.99988 | 
0.99981 | 


0.99973 
0.99963 
0.99952 
0.99940 
0.99927 


0.99913 
0.99897 
0.99880 
0.99862 
0.99843 


TEMPERATURE 











73 





DENSITY 


0.99823 
0.99802 
0.99780 
0.99756 
0.99732 


0.99707 
0.99567 
0.99406 
0.99224 
0.99024 
0.98807 
0.98573 
0.98324 
0.98059 
0.97781 
0.97489 
0.97183 
0.96865 
0.96534 
0.96182 
0.95838 





less, or 0.99718 gram. The following table, which gives the appar- 
ent weight of one liter of water weighed with brass weights in air 
and the corresponding correction to volume, is still more conven- 


jent. 


Apparent weight of one liter of water in air and corrections to be 
added to the apparent weight of one liter of water to find the true 
volume in cubic centimeters. 














eee as GRAMS CORRECTION 
1D. 998.05 z 
16° 997.90 2 
bye 997.74 2 
18° 997.56 | 2.49 ce. 
19° 997.38 2.65 ee 
20° 997.18 2.82 ee 
DABS 996.97 3.00 ce. 
AE 996.76 3.19 ee. 





TURE 


23° 
24° 
25° 
26° 
Poth 
28° 
29° 
30° 


TEMPERA- 








GRAMS CORRECTION 
996.53 3.40 ce 
996.29 3.61 ce 
996.04. 3.83 ee 
995.79 4.06 ee 
995.52 4.31 ee 
995.24 4.56 ee 
994.96 4.82 ee 
994.66 5.08 ec 














74 A TEXTBOOK OF CHEMISTRY 


Heat of Fusion and Vaporization. If heat is applied to a 
kilogram of ice at the freezing point, 0°, it will absorb 79 large 
calories in melting ; that is, if the same amount of heat is applied 
to a kilogram of ice and a kilogram of water, both of them at 0°, 
when the ice is melted and still at 0° the other water will be at 
a temperature of 79°. If a kilogram of water at the boiling 
point, 100°, is heated till it is all converted into steam at the 
same temperature and at atmospheric pressure it will absorb 
536 large calories, or enough to warm 5.36 kilograms of water 
from the freezing point to the boiling point. The steam formed 
will fill a space of about 1700 liters, while the liquid water fills 
only one liter. 

The heat which is absorbed in the melting of ice and vaporiza- 
tion of water was formerly called latent heat because it seems to 
disappear, but this expression is not as fre- 
quently used now as it was some years ago. 
The heat of vaporization varies with the tem- 
perature, being greater at lower and less at 
higher temperatures. 

Vapor Pressure of Water. If two dry 
tubes, about 800 mm. long, are filled with 
mercury and inverted, the mercury will fall 
until the weight of the mercury in the tubes 
is the same as the weight of a column of air 
of the same cross section and reaching to the 
top of the atmosphere. In other words, the 
mercury will stand at the same height as the 
mercury in a barometer. If a drop of water 
is introduced into one of the tubes, the mer- 

ey cury in the tube will fall and remain at a 
Fig. 30 position lower than that in the dry tube. If 
the tube containing the water is warmed, the 
mercury will fall further, if cooled, it will rise higher and for each 
temperature there will be a definite difference between the 
heights of the mercury in the two tubes. It is evident that this 
must be caused by the fact that a part of the water in the tube 





is converted into a vapor or gas and exerts a pressure on the 
mercury, partially balancing the pressure of the air. 
sure is called the pressure of water vapor, or sometimes, and less 
It is given for different tempera- 


VAPOR PRESSURE OF WATER 


correctly, the aqueous tension. 
tures in the following table: 


VAPOR PRESSURE OF ICE AND WATER 


This pres- 








PRESSURE IN 





ee ne a elec | Pn acces 
SNE 2.0 27° 26.5 100.5° (MENTE 
5° 3.0 28° 28.1 101.0° 787.6 
—/ 2° 3.9 29° 29.8 
==) |? 4.2 30° 31.6 
eas 4.6 31° 33.4 
Ge 4.9 32° 30.4 
2° 5.3 303° O74 
3° 5.7 34° 39.6 PRESSURE IN 
7, 6.1 35° 41.9 ATMOSPHERES 
5° 6.5 40° 50.0 
6° 7.0 50° 92.2 Pili 1.5 
fe 7.5 60° 149.2 120.6° 2 
8° 8.0 (Ab 233.8 127.8° 2.0 
ole 8.6 80° 309.5 133.9° 3 
103 9.2 90° 526.0 144.0° 4 
tie 9.8 95° 634.0 159.2° 6 
12° 10.5 96° 657.7 170.8° 8 
13° Lee Site 682.1 180.3° 10 
14° io 98° 707.3 188.4° 12 
15 Lea oe 733.2 195.5° 14 
16° 13.6 Ee 739.9 201.9° 16 
17" 14.5 99°2° 738.5 207.7° 18 
18° 15.4 99.3° 741.2 213.0° 20 
Uae 16.4 99.4° 743.9 224.7° 25 
20° 17.4 SS Rie 746.5 
ake 18.5 29.0% 749.2 
22° 17 DONS 7o1.9 
23° 20.9 00. Se 754.6 
24° 22.2 Dee 757.3 
25° 23.5 100.0° 760.0 
26° 25.0 100.1° 762.7 
100.2° 765.5 
100.3° 768.2 
100.4° 770.9 
































76 A TEXTBOOK OF CHEMISTRY 


Equilibrium. Solong as water remains in the liquid form in the 
barometer tube described in the last paragraph, the volume of 
the tube above will have no effect on the vapor pressure of the 
water. This relation, which is very important, will be clearer 
from Fig. 31. Suppose that the cylinder contains water in the 
bottom with vapor of water above it and that it is fitted with an 
air-tight piston. If the piston is raised, the vapor, which acts 
like any other gas in this respect, will immediately expand and 
fill this additional space, and the pressure will be momentarily 

lowered. Immediately, however, some of the water 
will evaporate, and this evaporation will continue till 
the pressure is the same as before, provided that the 
water is kept at the same temperature. If the piston 
is pressed down, the reverse operation will occur. The 
vapor will be momentarily compressed and the pressure 
increased, but some of the vapor will immediately con- 
dense to water and in this way the pressure will fall to 
its original value. When two or more forms of a sub- 
stance are related in this way, they are said to be in 
equilibrium, that is, they are so balanced against each 
other that any change in temperature or pressure will 
cause a partial conversion of one form into the other. 
It is not necessary to suppose that the same molecules 
Fig. 31 of water are always in the form of vapor in such a case. 
On the contrary, it seems more probable that some of 
the water particles are all of the time leaving the liquid and pass- 
ing into the vapor and that molecules of vapor are constantly 
leaving the vapor and passing back into the liquid. When the 
two are in equilibrium, just as many molecules must pass in one 
direction as in the other within a given time. 

Effect of Water Vapor on the Pressure of a Gas. If a small 
glass bulb filled with water (Fig. 32) is placed in a bottle filled 
with dry air and closed with a stopper bearing a manometer 
to show the pressure within the bottle, on breaking the bulb 
it will be seen that the pressure within the bottle increases and 
after some time the increase will be almost exactly equal to the 


PHASES fi 


vapor pressure of water at the temperature of the experiment. 
It is evident from this that as the water evaporates the vapor 
diffuses into the air above just as any other gas would do, and 
as it does so it adds its pressure to that of the air, in accordance 
with the law of partial pressures, that each gas in a mixture 
exerts the same pressure as though 
it were present alone (Dalton’s law 
of partial pressures). 

The experiment suggests the 
proper method of finding the vol- 
ume which a quantity of dry gas 
would fill under standard condi- 
tions, when the gas has been meas- 
ured in a moist condition. The 
actual pressure exerted by the gas 
is less than the apparent pressure 
by the pressure exerted by the 
vapor of water at the given tem- 
perature. For this reason in the 
experiment described on p. 67 
the direction was given to subtract 
from the reading of the barometer 
both the height of the mercury 
in the eudiometer and the pressure of water vapor, when 
the corrected volume of the moist hydrogen was to be 
found. 

Phases. Degrees of Freedom. Water may exist in the three 
forms of ice, water and vapor. These three forms of the same 
substance are called phases. As long as only one phase is present 
we may change either the temperature or pressure or both at 
will, and in order to know the condition of the phase we must 
know both the temperature and the pressure. The system is 
said to have two degrees of freedom and the system is called 
diwariant. 

When two phases are present, any change in the temperature 
will cause a corresponding change in the pressure, and as long as 


Fig. 32 


78 A TEXTBOOK OF CHEMISTRY 


the temperature is fixed the pressure cannot be changed without 
the disappearance of one of the phases. Or, as long as the pres- 
sure is fixed, the temperature cannot be changed. ‘Thus for 
water and vapor at a given temperature an increase of the volume 
does not decrease the pressure, but, instead of this, causes some 
of the water to change to vapor; and it is only when the liquid 
phase disappears that a further increase in the volume causes 
a decrease in the pressure. In the same way, if water and ice 
are present, an increase in the pressure will cause some of the 
ice to melt and the temperature will fall, and for every pressure 
there will be a corresponding temperature at which ice and 
water can exist together. To know the condition of such a 
system of two phases we need give only one factor. If we know 
the temperature, the pressure is fixed ; or if we know the pressure, 
the temperature is fixed by the properties of the substance. 
Such a system has only one degree of freedom and is called uni- 
variant. 

When the three phases, water, ice and vapor, are present, it 
is impossible to change either the temperature or the pressure 
without the disappearance of one of the phases. Such a system 
has no degree of freedom and is called invariant. 

The relations between temperature and pressure for the three 
phases of water may be seen clearly from the diagram (Fig. 33). 
If water and vapor are present, the relation between temperature 
and pressure is fixed by the line OA. For water and ice the re- 
lation is fixed by the line OC, from which it is apparent that an 
increase in the pressure lowers the melting point, though very 
slowly. The relation for ice and vapor is fixed by the line OB. 
The line OB’ gives the relation for vapor and supercooled water, 
the vapor pressure being slightly greater than that of ice at the 
same temperature. The line OB’ represents a condition of un- 
stable equilibrium, andif a little ice is introduced, the pressure will 
fall or the temperature will rise to the line OB, if all of the water 
freezes, or both pressure and temperature will rise to the 
point QO, which is the invariant point, called also the triple 
pownt. 


SOLUTIONS 79 


* The temperature of the system at the triple point will be 
+ 0.0073, since the 0° of our scale is determined by freezing water 
under atmospheric pressure and a pressure of one atmosphere 
lowers the freezing point 0.0073°. The vapor pressure of water 
at the freezing point is only 4.6 mm. 

Water as a Solvent. Solutes. If salt or sugar is placed in 
water, it quickly disappears and a homogeneous liquid is obtained, 
which we call a solution. Any substance which passes into solu- 
tion in this manner is called a solute. The liquid in which the 


PRESSURES 





-10 + 0.0073 +10 
TEMPERATURE 


Fig. 33 


solute dissolves is called a solvent. Solutes may be either solids, 
liquids or gases, and they vary very greatly in their degrees of 
solubility. Thus one liter of water will dissolve at 20° 670 grams 
of sugar or 358 grams of common salt, but it will dissolve only 
0.00153 gram of silver chloride. No satisfactory reason for such 
differences has been discovered, though many empirical relations 
between the composition of substances and their solubility are 
known. 

Some substances, as alcohol or sulfuric acid, dissolve in water 
in all proportions, but others will dissolve only up to a definite 
limit. When a solution can remain in contact with the solute 


80 A TEXTBOOK OF CHEMISTRY 


without taking up any more, it is said to be saturated. The solid, 
liquid or gaseous phase of the pure solute is then in equilibrium 
with the solution very much as vapor of water is in equilibrium 
with liquid water. 

The solubility of salts usually increases with the temperature, 
but there are some exceptions, and the rate of increase 
varies very greatly, as will be apparent from the accompany- 
ing diagram (Fig. 
34). 

If a warm. satu- 
rated solution of a 
salt which is more 
soluble in warm 


On 


than in cold water 
is allowed to cool 
out of contact with 
the solid phase, a 
supersaturated solu- 
tion may usually be 
obtained. The in- 
troduction of a little 
of the solid will 
start the separation 
of the solid phase, 
and after a short 
time the solution 
eo WL es mame 
0° 20° 40° 60° 80° 100° 120° 140° 160 

Temperature normal, saturated 
Fig. 34 condition. In a 
similar manner still 
water may be cooled below its freezing point or a vapor may 
be cooled or compressed below the point at which a part would 
ordinarily exist in the liquid form. Such a system is always 
unstable, somewhat after the analogy of a pyramid standing 
on its apex, and can only occur in the absence of the solid or 

liquid phase which should normally be present. 


SS 


ey 
Ee 
A 


IN 


t=) 

(ob) 

~ 

fas] 

5 

Ge 

° 

mM 

S| 

(a0 

te 

ao 

j=) 

(=) 

re 

g 

® 

tes i 
_— 

d PAR ei a 
mM 

D 

co} 

n 

g 

e 

1 Zz 
c= 

(o) 

> 

[a0] 

tH 

y=) 

fs] of 
oO 

(2) 

g 

(o) 

(@) 





SOLUTIONS Sl 


Chemical Activity in Solutions. Metathesis. If common salt, 
NaCl, and silver nitrate, AgNOs, are powdered and mixed to- 
gether, there will be no apparent action; but if each is dissolved 
in water and the solutions mixed, there will be formed immedi- 
ately a white precipitate of silver chloride; and if the solution 
is filtered from the precipitate and the filtrate evaporated, 
crystals of sodium nitrate may be obtained : 


NaCl + AgNO; = AgCl + NaNO; 
Sodium Chloride Silver Silver Sodium 
(common salt) Nitrate Chloride Nitrate 


Hundreds of illustrations of similar reactions which do not 
occur readily between the solid substances but which take place 
easily in solutions might be given. This reactivity of substances 
in solution is evidently in part because they are brought into inti- 
mate contact, since no combination can take place except be- 
tween substances which are touching each other. But this does 
not appear to be the only reason. In very many cases when 
clear, sharp-cut reactions occur, each compound separates, as 
here, into the metal and an acid radical. If we subject these 
same compounds in solution to the influence of an electrical 
current, the metal will travel toward the cathode through the 
solution while the acid radical will travel toward the anode. It 
seems, therefore, that solution in some way loosens the combina- 
tion between the ions so that they can very readily enter into 
new combinations. 

A reaction of the sort just considered is called a double decom- 
position or metathesis. 

Hydrates, Deliquescence, Efflorescence. Many salts when 
they separate in crystals from solution do so in combination with 
a definite quantity of water. Thus crystals of copper sulfate 
have the composition CuSO..5 H.O; crystals of sodium sulfate, 
the composition NazSO4.10 H,O. Such compounds usually 
decompose rather easily into water and the anhydrous salt, and 
the water is spoken of as water of hydration, and the compounds 


82 A TEXTBOOK OF CHEMISTRY 


are called hydrates. If the hydrate is placed above the mercury 
in a barometer tube it will decompose, giving off water vapor till a 
definite vapor pressure is reached. This vapor pressure will vary 
greatly for different hydrates and will increase with the tem- 
perature as the vapor pressure of water does. Thus the vapor 
pressure of the hydrate CuSO,.5 H.O is 12.5 mm. at 30°, the 
vapor pressure of CaCl,.H2O is only 3.1 mm., the vapor pressure 
of NagSO4.10 H.O is 25.3 mm. and the vapor pressure of pure 
water is 31.6 mm. If an anhydrous salt like calcium chloride is 
exposed to air containing enough moisture so that the pressure 
of the water vapor in it exceeds: 3.1 mm. at 30°, water will be 
absorbed and the hydrate will be formed. In this case even a 
concentrated solution of calcium chloride has so low a vapor 
pressure that it is exceeded by that of the moisture in ordinary 
air. Accordingly calcium chloride when exposed to the air 
absorbs moisture and finally dissolves in the water absorbed. 
Such a salt is said to be deliquescent. 

On the other hand, if the hydrate, Na2SO..10 H,O, is ex- 
posed to air in which the vapor pressure of the water which it 
contains is less than 25.3 mm. at 30° the salt will decompose and 
lose water to the air. As it does so it will fall to a fine powder or 
flour. Salts of this type are called efflorescent. 

Natural Waters. The water which is found in nature is never 
pure, the purest being rain water falling in the open country 
after it has been raining for some time, or water obtained by 
melting the ice from a pure, fresh-water lake. Even such water 

1 The term water of crystallization, which is still used by many 
authors, is not as appropriate. During the first half of the nine- 
teenth century sodium hydroxide, NaOH, and calcium hydroxide, 
Ca(OH)», were called sodium hydrate and calcium hydrate, and their 
formulas were written in a form which with modern atomic weights 
would be Na:O.H;0 and CaO.H.0. This older use of the word 
hydrate, which still clings to the literature of pharmacy, has inter- 
fered somewhat with the introduction of the word in the sense in 
which it is here defined. 

It is fair to say, too, that the distinction between hydrates and 
hydroxides is more or less arbitrary, as some hydroxides lose water 
more easily than some hydrates, and in many cases the water of 


hydration cannot be removed without decomposition of the rest 
of the salt. 


NATURAL WATERS 83 


contains air in solution and a little carbonic acid from the carbon 
dioxide of the air. On falling upon the ground rain water begins 
at once to take up various substances, partly in suspension, partly 
in solution. Calcium or magnesium carbonate and calcium 
sulfate cause the water to become hard (pp. 310, 311) and injure 
it seriously for use in steam boilers or in laundries. If the water 
is mixed with sewage, it frequently becomes contaminated with 
bacteria which produce disease. The diseases of typhoid fever 
and of cholera, especially, are frequently transmitted in this way. 
During an epidemic of cholera in the cities of Hamburg and 
Altona, Germany, the people in the houses on one side of a certain 
street used a contaminated water which caused very many cases 
of the disease. Across the street, water from the same source, 
but after passing a system of public filters, was used and there 
were very few cases. In Chicago before the opening of the 
drainage canal there were 170 deaths from typhoid fever per 
year for each 100,000 people. After the drainage canal carried 
away the sewage which had formerly gone into Lake Michigan 
and contaminated the water supply of the city, the death rate 
from typhoid fell to 16 per 100,000. 

Purification of Water. Water may be purified most com- 
pletely by distillation, though special precautions are required 
to get rid of ammonia, carbon dioxide and other volatile impuri- 
ties. The bacteria in water may be killed almost completely 
by boiling the water for a short time, and this should always be 
done when it is necessary to use a suspected water for drinking 
or the preparation of food. The bacteria may also be almost 
completely removed by filtration, either on a large scale on beds 
of sand, or through filters of fine-grained stone or unglazed por- 
celain. Charcoal filters, which were formerly used, are rarely 
effective. Waters may also be sterilized by treatment with 
ozone, with ultra-violet light, or with bleaching powder. 

Hydrogen Peroxide. A second compound of hydrogen with 
oxygen, called hydrogen peroxide and having the formula H.O2, 
can be prepared by the action of acids on peroxides of univalent 
or bivalent metals : 


UNIVERSITY OF ILLINOIS 


LIBRARY-CHEMISTHI 


84 A TEXTBOOK OF CHEMISTRY 


Na2Oz -+ 2 HCl 2 NaCl -+- H,0, 


Sodium Peroxide . Hydrogen Peroxide 
BaO, + H,SO, — BaSO, j 4+ H,0, 
Barium Peroxide Barium Sulfate 


Barium sulfate is almost insoluble in water, and if the barium 
peroxide and sulfuric acid are used in equiv- 
alent amounts a solution may be obtained 
which contains practically nothing except 
hydrogen peroxide and water. Such a solu- 
tion may be concentrated in a vacuum des- 
iccator (Fig. 35) over sulfuric acid, the 
vapor pressure of water being much greater 
than that of hydrogen peroxide, or it may 
be distilled under the low pressure obtained 
by means of a good air pump or filter pump. 
Water will distil away first, and finally the 
hydrogen peroxide will distil, almost pure. 

Hydrogen peroxide is a heavy liquid which 
decomposes slowly at ordinary temperatures 
into water and oxygen: 



































H.C, = H.O + O 


This reaction is accompanied by the evolution of 46,200 cal- 
ories of heat (by the decomposition of 68 grams of hydrogen 
peroxide), and, as is usual with reactions evolving heat and lib- 
erating a gas, may become explosive. The reaction is catalyzed 
by many substances, especially by finely divided platinum and 
other metals, and it is quite dangerous to bring concentrated 
solutions of hydrogen peroxide into contact with organic sub- 
stances. 

Hydrogen peroxide is a powerful oxidizing agent, giving oxygen 
readily to many substances. This property, also, is closely re- 
lated with the fact that heat is evolved when it decomposes. 

If a solution of hydrogen peroxide is added to the black pre- 


HYDROGEN PEROXIDE 85 


cipitate of lead sulfide suspended in water, the substance is oxi- 
dized to white lead sulfate : 


PbS + 4 HO. = PbSO,. + 4 H.O 


Hydrogen peroxide also acts on silver oxide and on a good many 
other substances as a reducing agent, but in all such cases oxy- 
gen is liberated. We may explain this by considering that 
oxygen gas is a compound of oxygen with itself ! and that the oxy- 
gen of the hydrogen peroxide oxidizes the oxygen of the silver 
oxide to free oxygen. Or it may be that the oxygen of the 
silver oxide oxidizes the hydrogen of the hydrogen peroxide to 
water, leaving the two oxygen atoms of the peroxide combined 
together as free oxygen, O2?: 


H.0, at. AgsO ee H.O + O. 4. 2Ag 


The reduction of potassium permanganate, KMnQu,, in 
acid solution is frequently used for the quantitative deter- 
mination of hydrogen peroxide. The equation representing 
this reaction is most easily written in two parts, as indicated 
below. The first part of the reaction is based on the facts that 
potassium is univalent and manganese bivalent when combined 
with acid radicals and that, for this reason, when two molecules 
of potassium permanganate are acted on by sulfuric acid in the 
presence of some substance which can take up oxygen, five atoms 
of oxygen become available for the oxidation of that other sub- 
stance. These five atoms of oxygen are inclosed in brackets 
to indicate that the decomposition does not take place except in 
the presence of something with which this oxygen can combine: 


2 KMnO, + 3 H2SO4 = K2SO4 + 2 MnSO, + 3 H20 + (5 O) 
(50) + 5 H.02 = 5 H20 + 5 Op 
2 KMnO, + 3 H2S0, + 5 H2O2 
= K,S0,4 + 2 MnSO, + 8 H,O + 5 Op 


1 Other reasons for such a view will be given later (p. 93). 
2 The fact that the reaction 
H,0, = H,.O+ 0 
is monomolecular points strongly to the latter explanation. 


86 A TEXTBOOK OF CHEMISTRY 


The last equation is obtained by combining the other two 
algebraically, eliminating the five atoms of oxygen which appear 
on opposite sides. 

Properties and Uses of Hydrogen Peroxide. Hydrogen 
peroxide has a specific gravity of 1.4584 at 0°. It boils at 69.2° 
under a pressure of 26 mm. or at 84°-85° under a pressure of 
68 mm. A dilute solution is fairly stable when pure, the stabil- 
ity being increased by the presence of a small amount of a mineral 
acid. The stability is decreased by alkalies and by many other 
substances, and the decomposition is also hastened by exposure 
to light. 

Hydrogen peroxide is used in medicine as a bactericide and for 
the diagnosis of pus, which causes its rapid decomposition with 
evolution of oxygen. The medicinal solution commonly used 
is known as a 10-volume solution, meaning that it evolves ten 
times its volume of oxygen when it decomposes, or 20 times its 
volume when treated with an oxidizing agent, one half of the 
oxygen coming from the latter, as explained above. 

Hydrogen peroxide is also used to bleach hair, silk and wool, 
being much more suitable than chlorine or hypochlorites (p. 127) 
for this purpose. A solution suitable for this purpose may be 
obtained by dissolving sodium peroxide in cold water and adding 
dilute sulfuric acid. 

* Tests for Hydrogen Peroxide. With potassium dichromate 
in an acid solution hydrogen peroxide gives a beautiful blue 
compound which is soluble in ether. The composition of the 
compound is not positively known. Its formation may be 
used either as a test for hydrogen peroxide or as a test for chro- 
mium. Another valuable test is the yellow color given with 
solutions containing titanium. . 

*Structure of Hydrogen Peroxide. Two formulas have 
been proposed to represent the structure of hydrogen peroxide: 


H 
HOO Hend »0=0 
H 


LAW OF MULTIPLE PROPORTION 87 


The first represents both oxygen atoms as bivalent, the second 
represents one of the oxygen atoms as quadrivalent. The evi- 
dence in favor of the first formula is: 

1. There are a number of reactions in which hydrogen peroxide 
seems to be formed by the reduction of oxygen, which has the 
formula Oz: (p. 93) : 


O=O0O + H—H = H—O—O—H 


2. Hydrogen peroxide does not seem to be formed by the oxi- 
dation of water, as it should be if the second formula were true. 

3. When the two hydrogen atoms of hydrogen peroxide are 
replaced by the ethyl group, C2Hs, diethyl peroxide, (C.H;)2O., 
is formed and this gives ethyl alcohol, C.H;—O—H, by re- 
duction. ‘This is easily explained by the first formula: 


C.H;—O—O—C2H; + H, —— Pe C,H;—O—H 


If diethyl peroxide had a formula corresponding to the 
second one for hydrogen peroxide, it should give ethyl ether, 
C.H;—O—C2H; by reduction (Baeyer and Villiger, Ber. 33, 
3387 (1900)) : 


Law of Multiple Proportion. In water and hydrogen peroxide 
the same elements, hydrogen and oxygen, combine to form two 
different compounds. In water one part of hydrogen combines 
with eight parts of oxygen (or 2 with 16), while in hydrogen per- 
oxide one part of hydrogen combines with sixteen parts of oxy- 
gen (or 2 with 32). Nitrogen and oxygen form a whole series 
of compounds represented by the formulas and composition : 


N:O 
Nitrous oxide, NsO_ - 28:16 
Nitric oxide, NO 14:16 


Nitrogen trioxide, N20; 28:48 
Nitrogen tetroxide, NeO, 28: 64 
Nitrogen pentoxide, N2O; 28:80 


88 A TEXTBOOK OF CHEMISTRY 


In every such case, if we consider some fixed amount of one of 
the elements (it makes no difference which one) the amounts of the 
other combining with this fixed amount will bear a simple ratio to 
each other. This is known as the Law of multiple proportions. 
The discovery of this law led Dalton to propose the atomic 
theory of the constitution of matter. A little consideration of 
the law shows that it follows, necessarily, from the law of com- 
bining weights (p. 13), and that the law of combining weights is 
more comprehensive and important. 

Two compounds of iron contain: 


Ferrous oxide, 77.73 per cent of iron 
22.27 per cent of oxygen 
100. 

Ferric oxide, 69.94 per cent of iron 
30.06 per cent of oxygen 
100. 








Show that these proportions are in accordance with the law 
of multiple proportion. This illustration helps us to understand 
why the law was not discovered before the nineteenth century. 


ae 


CHAPTER VI 


AVOGADRO’S LAW. SELECTION OF ATOMIC WEIGHTS. 
OZONE 


Gay Lussac’s Law of Combining Volumes. It has been shown 
that hydrogen and oxygen unite very nearly in the proportion 
of two to one by volume and that the volume of steam formed is 
very nearly the same as the volume of hydrogen which it contains. 
A study of many other gaseous elements and their compounds 
has shown that in every case there 1s a simple ratio between the 
volumes of gases which combine with each other and also between 
those volumes and the volume of the product, if that 1s a gas. This 
is known as Gay Lussac’s law of combining volumes, and it is 
true of all elements or compounds which can be converted into 
gases without decomposition, as well as of substances which are 
gases at ordinary temperatures. It has been shown to be true 
for thousands of compounds. The law may be illustrated by the 
following diagrams : 





2 grams 18 grams 

Hydrogen 32 grams Steam 
+ Oxygen — 

2 grams 18 grams 

Hydrogen 1 vol. Steam 





2 vols. 2 vols. 


2 grams 71 grams 


Hydrogen ah Chlorine 





1 vol. lvol. 89 


90 A TEXTBOOK OF CHEMISTRY 


2 grams 
Hydrogen 
2 grams 28 grams 17 grams | 17 grams 
Hydrogen =F Nitrogen Ammonia | Ammonia 





1 vol. 2 vols. 
2 grams 


Hydrogen 





3 vols. 


271 grams 


200 grams 71. grams 


Mercuric 


Mercury - Chlorine Chloride 


2 grams 2 grams 


Hydrogen | Hydrogen 


34 grams | 34 grams 
Phosphine | Phosphine 
2 grams 


Hydrogen | Hydrogen 
34 grams | 34 grams 





2 grams 2 grams Phosphine Phosphine 





Hydrogen | Hydrogen 





The volumes have been chosen for these illustrations in such a 
way that the unit volume always contains one gram, or a whole 
number of grams, of hydrogen. In every case the weight of this 
unit volume bears a simple relation to the atomic weights of the ele- 
ments which it contains. This volume of any gaseous element or 
compound always contains one gram atom or a whole number of 
gram atoms of each element contained in the gas. This is not 
accidental, but follows of necessity from the two laws: (1) that 


AVOGADRO’S LAW 91 


the composition of every compound can be expressed by multi- 
ples of the atomic weights of the elements which compose it ; 
and (2) that there is always a simple ratio between the volumes 
of gases which combine and also between those volumes and the 
volume of the product, if that is a gas. 

A further examination of the illustrations shows that the unit 
volume of each compound contains one gram molecule of the 
compound. ‘This result does not follow of necessity from the 
two laws just given. It depends on the values which we give to 
the atomic weights. Thus if we were to call the atomic weight 
of oxygen 8 and the formula of water HO, as was done by Dalton, 
the gram molecular weight of water would be 9 and the unit 
volume which has been chosen would contain two gram mole- 
cules of water. As has been stated, we have chosen as the unit 
volume for our illustration that volume which, in several com- 
pounds, contains one gram of hydrogen, of course because hy- 
drogen is our (approximate) unit for atomic weights. No com- 
pound of oxygen is known which contains less than 16 grams in 
this unit volume. 

Avogadro’s Law. Theselectionof 16 instead of 8 as the atomic 
weight of oxygen is based on a hypothesis proposed by Avogadro, 
an Italian chemist, in 1811. This hypothesis is that all gases, 
under the same conditions of temperature and pressure, contain 
equal numbers of molecules in equal volumes. Some of the many 
facts which support this hypothesis so fully that it may now be 
considered as an established law are: 

1. While solids and liquids vary greatly in their rate of expan- 
sion or contraction for changes of pressure or of temperature, all 
gases expand and contract alike. This points very strongly to 
a similarity in their structure. 

2. When the law of combining weights is combined with the 
law of combining volumes it follows, of necessity, that there must 
be a simple ratio between the numbers of molecules in equal 
volumes of different gases. But if the ratio is one of simple 
whole numbers, it seems highly improbable that it is not one of 
equality. Thus it would seem very improbable that the number 


92 A TEXTBOOK OF CHEMISTRY 


of molecules in a given volume of steam is exactly twice the 
number in the same volume of hydrochloric acid, as we should 
have to suppose if the molecular weight of water is 9 while that 
of hydrochloric acid is 36.5. 

3. The kinetic theory of gases, which accounts so well for their 
properties, leads directly to Avogadro’s law on the basis of the 
laws of the collision of elastic bodies (p. 58). 

4. The atomic weights selected on the basis of Avogadro’s 
law have made possible the classification of the elements in 
groups according to their atomic weights. This classification 
is known as the Periodic System (p. 132) and furnishes very 
strong, independent evidence that the atomic weights selected 
are in reality the true relative weights of the atoms. 

Selection of an Atomic Weight. According to Avogadro’s law, 
equal volumes of different gases, under the same conditions of 
temperature and pressure, must have weights proportional to 
the weights of the molecules of the gases. If the molecule of one 
gas is twice as heavy as that of another, one liter of the first gas 
must weigh twice as much as a liter of the second. Accordingly 
if we can take as our unit volume the volume filled by a gram 
molecule of some compound which contains only one atom of 
hydrogen in its molecule, this unit volume will contain one gram 
molecule of every other gaseous element or compound. We can- 
not, of course, look at the molecules of different compounds to 
discover which one contains only a single atom of hydrogen; but 
if we weigh the same volume of hydrochloric acid, steam and 
ammonia, we find that in a given volume of steam there is twice 
as much hydrogen as there is in the same volume of hydrochloric 
acid, and thatin the same volume of ammonia there is three times 
as much hydrogen. As no compound has ever been found which 
contains less hydrogen than hydrochloric acid does in the unit 
volume, it seems pretty safe to conclude that there is only one 
atom of hydrogen in a molecule of this compound. As one gram 
of hydrogen combines with 35.5 grams of chlorine,! the volume 


1 Approximate values are used for convenience, as always. The 
true values are 1.0078 grams of hydrogen for 35.46 grams of chlorine. 


SELECTION OF ATOMIC WEIGHTS 93 


filled by 36.5 grams of hydrochloric acid will contain one gram 
molecule, and it must contain one gram molecule of every other 
gas. 

In the discussion above it is pointed out that those compounds 
of hydrogen which contain the smallest amount of hydrogen in 
the unit volume probably contain only one atom of hydrogen 
in their molecules. Reasoning in the same way, we may find 
for each element those compounds which contain the smallest 
amount of the element in the unit volume, and it is probable that 
these compounds contain only one atom of the element in the 
molecule. Thus steam contains 16 grams of oxygen in the unit 
volume; and as no compounds of oxygen containing a smaller 
amount is known, we conclude that there is only one atom of 
oxygen in a molecule of steam and that the atomic weight of 
oxygen is 16. The atomic weights of the other elements which 
form gaseous compounds, or compounds which can be converted 
into gases without decomposition, have been selected in the same 
way. 

Molecules of the Elements. It was natural for Dalton when 
he proposed the atomic theory to think of the atom as the small- 
est particle of an element in the free state, and it did not occur 
to him that atoms of the same kind could combine. A refer- 
ence to the diagrams, however, shows that there is twice as much 
oxygen in the unit volume of oxygen gas as in the unit volume 
of steam. According to Avogadro’s law it follows that oxygen 
gas contains two atoms of oxygen in each molecule. We may 
reach the same conclusion by another process. If we let each 
square below represent 1000 molecules, it is clear that 2000 
molecules of steam are formed from 1000 molecules of oxygen, 
and, as each molecule of steam must contain at least one atom 


1000 | 1000 
M 


g Molecul 
olecules 1000 olecules 


H H 
2 — Molecules os 20 
1000 


1000 
Molecules 
H:O 


O» 





Molecules 
He 





94 A TEXTBOOK OF CHEMISTRY 


of oxygen, the 1000 molecules of oxygen must consist of 2000 
atoms, or each molecule contains two atoms. 

It will be seen from the diagrams that four atoms combine 
to form a molecule of phosphorus vapor, while a molecule of 
mercury vapor contains only a single atom. In the latter case 
atom and molecule are identical. In general the atoms of 
nonmetallic elements combine to form molecules of the element 
in the free state, but the atoms of the metals show little tendency 
to combine in this way. For a probable explanation of this 
remarkable difference see J. J. Thomson, The Corpuscular 
Theory of Matter, p. 120; and H. N. McCoy, J. Am. Chem. 
MOCO pono: 

Gram Molecular Volume. The unit volume, which will con- 
tain one gram molecule of any gas, is best calculated from the 
weight of a liter of oxygen, since oxygen is the basis for atomic 
weights (p. 68). One gram molecule of oxygen, Os, contains 
32 grams, and, as a liter of oxygen weighs 1.429 grams, the gram 
5 = 22.4 liters at 0° 
and 760 mm. At the same temperature and pressure one 
gram molecule of any other gas will fill the same volume, 22.4 
liters. 

This statement and the law of Avogadro are subject to limita- 
tions similar to those which apply to the laws of Boyle and 
Charles. Just as most gases, and especially those which are 
easily liquefied, contract too much when compressed from one 
to two atmospheres pressure and also contract too much when 
cooled from 100° to 0°, so almost all gases are heavier than they 
should be in accordance with the law of Avogadro. As the 
volume increases under diminished pressure, however, gases ap- 
proach the condition of an “ideal ”’ gas, and at low pressures the 
densities of gases agree very closely indeed with the law. The 
amount of the deviation from the law at atmospheric pressure 
and the agreement under low pressures will be obvious from the 
following table : 





molecular volume for oxygen will be 


DENSITY OF GASES 95 


Density oF GASES! 























WEIGHT OF | DENSITY AT Minne Gas 
Newz Fonaosa) Owe Large | az O ano | Parspume | wiA@, 
eee 2, a”. | Os 1.429 32.00 32.00 SP, 
Hydrogen . .| Hp 0.08987 1.997 2.01 2.016 
Nitrogen . . .| Ne 1.2507 28.02 28.01 28.02 
Carbon monoxide} CO 1.2504 28.01 28.00 28.00 
Nitric oxide . .| NO 1.3402 30.02 30.01 30.01 
revokes: ot, ALP 1.7808 39.89 39.88 39.88 
Carbon dioxide .| CQO, | 1.9768 44,28 44.01 44.00 
Nitrous oxide .| NO | 1.9777 44.30 44.03 44.02 
Hydrochloric acid) HCl | 1.6398 36.73 36.47 36.47 
Ammonia . .| NH3j| 0.7708 TR20, 17.01 17.03 
Sulfur dioxide .| SO, | 2.9266 65.56 64.07 64.07 
Air. ....| — 1.2928 28.96 — ae 

















The table shows that at low pressures the deviations from Avo- 
gadro’s law searcely exceed the experimental errors of the deter- 
minations. If the pressure and volume of a gas are corrected for 
the volumes occupied by the molecules and for their attractions for 
each other, corrections which can be determined experimentally 
(van der Waals), the law of Avogadro also becomes almost rigor- 
ously exact. 


The weight of 22.4 liters of air furnishes a very simple method 
of calculating the density of any gas as compared with air. 
For approximate estimates the weight may be taken as 29 grams, 
and the molecular weight of any compound divided by 29 will 
be approximately its density as compared with that of air. 
Thus it will be seen that hydrogen is 145 times lighter than air, 
while oxygen is 1.1 times as heavy and carbon dioxide one and 
one half times as heavy. 

* Number of Molecules in one Cubic Centimeter of a Gas. 
Avogadro’s law was established with a high degree of probability 
before any means of estimating the number of molecules in a 


1 Guye, J. Am. Chem. Soe. 30, 155 (1908). 


a 


96 A TEXTBOOK OF CHEMISTRY 


given volume of a gas was discovered. It is of some interest 
to know, however, that several different methods of estimating 
this number are now known and that the results obtained by 
different methods are in fair agreement. ‘Two of these methods 
may be given here, in outline. 

It has been shown that the element radium slowly decomposes 
and that as it does so it shoots out atoms of helium with a tremen- 
dous velocity. The volume of helium given off by a gram of 
radium has been measured and is 0.46 cubic millimeter in a day. 
When an atom of helium shot out by radium hits a screen of zine 
sulfide, it produces a flash of light; and as the helium atoms are 
sent out in all directions equally, by placing a screen of zinc 
sulfide back of a small opening which is at a known distance 
from a weighed amount of radium and counting the flashes, it 
is possible to estimate the number of atoms of helium which pass 
the opening in a given time and so the number of atoms in a 
cubic centimeter of the gas. Millikan has estimated the num- 
ber of molecules in one cubic centimeter of helium, in this 
way, as 2.71 X 10! (or 27,100,000,000,000,000,000). (See also 
Rutherford, Report of the Winnipeg Meeting of the British 
Association, 1909, p. 377.) 

When very minute particles suspended in water are observed 
with a microscope, it is seen that they are never at rest, but move 
about constantly in a wholly irregular manner. This was first 
noticed by the English botanist Brown in 1827 and is called the 
Brownian movement. It has long been considered as an evi- 
dence that water and other liquids are composed of molecules 
which are in rapid motion. Perrin has succeeded in showing 
(1909) that in an emulsion of gamboge the minute particles are 
distributed as they should be if they are considered as very large 
molecules, and he has connected this distribution with the kinetic 
theory in such a manner as to furnish an estimate of the number 
of molecules in a cubic centimeter of a gas, which he gives as 
3.15 X 10%. A number of other methods of estimating the 
same quantity give results of the same general value. Ags the 
different methods are quite independent of each other, we can 


OZONE 97 


have considerable confidence that the values are approximately 
correct.} 

Allotropic Forms. Ozone. When oxygen is subjected to the 
action of electrical waves obtained by connecting the tinfoil at 
A and D (Fig. 36), with the poles of an induction coil, from 5 
to 8 per cent of the gas is converted into ozone. If the mixture 








XQ GQ 99.54 





Y) (J, 


ViiiiiLlllllililllllllilttlttittlttltttltlttlittttiililitiiéititltllliltlttltllizZeées 
Fig. 36 


of oxygen and ozone is partly liquefied by passing it into a flask 
immersed in liquid air, a dark blue liquid is obtained. Oxygen 
boils at — 182.5°, while ozone boils at — 119°, and a mixture of 
oxygen and ozone containing 84 per cent of ozone has been ob- 
tained by allowing the oxygen to boil away from such a mixture. 
(Ladenburg, Ber. 31, 2508, 2830 (1898).) 

Ozone has a strong odor, which is noticed in the neighborhood 
of electrical machines. The weight of the gram molecular 
volume (22.4 liters) of the gas is 48 grams. Hence the formula 
is.O3. Ozone is evidently formed in accordance with the equa- 


tion: 
3 O, wie / O; 


The reaction is reversible and the ozone formed is unstable, 
decomposing slowly at ordinary temperatures, rapidly and 
completely at 250°-300°. The decomposition is accompanied by 
considerable evolution of heat, and so liquid ozone may easily 
give violent explosions. One gram molecule, 48 grams, gives 
out 29,400 calories in decomposing, and the properties of ozone 
are intimately connected with the fact that it contains so much 
more chemical energy than ordinary oxygen. It is a vigorous 


98 A TEXTBOOK OF CHEMISTRY 


oxidizing agent and attacks metallic silver and many other sub- 
stances which are not affected by ordinary oxygen. 

Ozone may be detected by its action on moist potassium 
iodide starch paper : 


The liberated iodine gives a deep blue color with the starch. 
A small amount of ozone is formed when a piece of clean phos- 
phorus, half covered with water, is allowed to stand for a short 
time in a bottle filled with air. The ozone may be detected by 
the odor and by the potassium iodide starch paper. 

Ozone has been used to a limited extent for sterilizing drinking 
water, but this use has not proved very successful. Bleaching 
powder is cheaper and much more efficient. Ozone is, however, 
a powerful germicide, and it may be that the ozone generated in 
thunderstorms plays a beneficent part in nature. 

In ordinary oxygen and ozone we have two different forms of 
the same element, and it is evident that in this case we may ob- 
tain another substance from an element without adding any- 
thing toit. The formulas of the two forms of oxygen give us a 
partial explanation of the difference between the two substances. 
Both forms are really compounds in which oxygen is combined 
with itself. We can represent this graphically by the formulas : 


O 
O=O and “Y. or O=O=O0 
Oso 


The last formula represents one of the oxygen atoms as 
quadrivalent. 

Several elements beside oxygen exist in two or more forms. 
Such forms are called allotropic. 


OZONE 99 


EXERCISES 


1. What will be the volume of a gram molecule of oxygen at 20° 
and 750 mm. ? 


2. How many grams of sodium peroxide will be required to give 
22.4 liters of oxygen under standard conditions ? 


3. How many grams of potassium chlorate will be required to give 
22.4 liters of oxygen ? 


4. How many liters of carbon dioxide will be formed by burning 6 
grams of carbon ? 


5. How many liters of carbon monoxide? How many liters of 
oxygen will be required for the reaction in each case ? 


6. How many grams of iron will be required to give 22.4 liters of 
hydrogen, if the iron is dissolved in hydrochloric acid? How many 
grams, if dissolved in sulfuric acid? How many grams, if used to 
decompose steam ? 


=# How much heat could be obtained, if two grams hydrogen could 
be burned in ozone ? 


8. How many grams of sodium are required to give 22.4 liters of 
hydrogen ? 


9. How many milligrams of aluminium are required to give 22.4 ce. 
of hydrogen ? 


CHAPTER VII 


CHLORINE 
Sympou, Cl. Atomic Weieuat, 35.46. Formuua, Clo. 


Occurrence of Chlorine. Chlorine is not found free in nature, 
chiefly because of its strong affinity for almost every other ele- 
ment and especially for the metals. Its most important com- 
pound is common salt, sodium chloride, NaCl. This is found in 
large amounts in sea water, in strong brines from artesian wells 
in very many places and in enormous beds of rock salt, some- 
times a hundred feet in thickness. Calcium chloride and mag- 
nesium chloride are also found in sea water and in many of the 
brines. Silver chloride is sometimes an important ore of silver. 

Preparation of Chlorine. 1. By Electrolysis of Sodium Chlo- 
ride. If an electric current is passed through a solution of 
common salt, NaCl, using a carbon anode and mercury as a 
cathode, the negative chloride ions, Cl-, will be discharged at the 
anode and the chlorine will be evolved as a gas. The sodium 
will dissolve in the mercury cathode, and by appropriate means it 
may be caused to react with water, giving sodium hydroxide and 
hydrogen. ‘This process is now used extensively in the manu- 
facture of chlorine and caustic soda or sodium hydroxide. It 
is known as the Castner-Kellner process and will be considered 
further in connection with sodium hydroxide (p. 402). 

2. Preparation by Oxidation of Hydrochloric Acid. In the 
preparation of oxygen, compounds (mercuric oxide or potassium 
chlorate) are selected which decompose with liberation of the 
gas when heated. In the preparation of hydrogen, compounds 
(water or hydrochloric or sulfuric acid) are selected from which 
the hydrogen can be displaced by another element. Both 


methods may be used for the preparation of chlorine. 
100 


CHLORINE 101 


In all of the methods practically used except electrolysis, the 
chlorine of hydrochloric acid is displaced by oxygen, one atom 
of the bivalent oxygen displacing two atoms of chlorine : 


H—Cl DADS 


In several methods compounds of manganese are used, and 
these methods depend upon the fact that while manganese forms 
compounds with oxygen in which it is quadrivalent, or in which 
it even shows a higher valence, the compounds of the element 
with chlorine which contain more than two atoms of chlorine 

for one of manganese are very unstable. The compounds with 


oxygen are manganese dioxide, Mn€ , and potassium perman- 


ganate, K—O—Mn€0. The stable compound with chlorine 
/CI, 
mt] 

If a concentrated solution of hydrochloric acid is poured on 
manganese dioxide, there is formed, at first, a dark brown solu- 
tion which probably contains some manganese tetrachloride : 


is Mn 


MnO, + Heh = wihinGl, 2 Fs 2: HO 
Manganese Hydrochloric Manganese 
Dioxide Acid Tetrachloride 


This equation represents the reaction as a metathesis in which 
two compounds each separate into two parts and then one part 
of each combines with one part of the other compound. A very 
large majority of chemical reactions belong to this class, and 
the principle is very useful in writing and understanding chemical 
equations. In the present case the manganese tetrachloride 
is so unstable that only indirect evidence of its existence has 
been obtained. When the solution is warmed gently, it is de- 
composed completely into manganous chloride and chlorine : 


MnCl, = MnCl. +. Cl, 


102 A TEXTBOOK OF CHEMISTRY 


The two reactions, which go on simultaneously, may be ex- 
pressed in one equation, thus : 


MnO, + 4 HCl = MnCl, + Cl, + 2 HO 


If concentrated hydrochloric acid is al- 
lowed to drop on potassium permanganate, 
KMnO,, which is a powerful oxidizing 
agent, the hydrochloric acid is oxidized 
and chlorine is liberated. A part of the 
chlorine, of course, remains combined 
with the potassium and manganese. In 
writing the equation we notice that we 
must have 8 molecules of hydrochloric 
acid to furnish the hydrogen to combine 
with the 4 atoms of oxygen in one mole- 
cule of the potassium permanganate, and 
we write, at first: 


KMnO, + 8 HCl = 
KCl + MnCl, + 4 H;,0 + 5 Cl 


But this form of the equation gives an 
odd number of atoms of chlorine, and as 
the chlorine is liberated in the molecular 
form it is necessary to double the equation to represent the 
substances actually formed : 


2 KMnO, + 16 HCl = 2 KC] + 2 MnCl. + 8 H.O + 5 Cl, 





3. Preparation of Chlorine by the Deacon Process. In the 
methods described in the last paragraph for the preparation of 
chlorine the element is liberated from hydrochloric acid by 
means of an expensive oxidizing agent. Henry Deacon of Eng- 
land devised a process a good many years ago by which the ex- 
pensive oxidizing agent is replaced by the oxygen of the air. 
If a mixture of hydrochloric acid and oxygen, or air, is heated, 
the reversible reaction represented by the equation : 


4 HCP 0722 HO 2. Ch 


CHLORINE 103 


takes place, but at a temperature of 350° to 400° the reaction is 
too slow to be commercially possible, as time is a very important 
element in all technical processes. The reaction might be has- 
tened by using a higher temperature, but at a high temperature 
the equilibrium between the four substances is displaced in such 
a way that less chlorine is formed and more of the hydrochloric 
acid passes through the heated apparatus unchanged (p. 110). 
By the use of a catalyzer, however, the reaction at lower tem- 
peratures can be hastened so far as to become technically pos- 
sible. 

Several different catalyzers may be used, but Deacon found 
that copper chloride, in the form obtained by saturating pumice 
with a solution of the salt and drying, is most suitable. With 
this catalyzer the reaction is sufficiently rapid so that, at 345°, 
80 per cent of the hydrochloric acid can be oxidized to chlorine. 
The process does not, however, compete successfully with other 
processes for the manufacture of chlorine. 

* 4. The Weldon Process for Chlorine. The oxygen of the 
air is also used, indirectly, as the oxidizing agent for the hydro- 
chloric acid in the Weldon process. When calcium hydroxide 
(in the form of milk of lime) is added to the solution of manganese 
chloride obtained in the preparation of chlorine by means of 
manganese dioxide, a precipitate of manganese hydroxide is 
formed : 

Gl ve H 


OH Cl 
Mint + Ca ie + Gra 
Cl 


= Mn 
Norn ancl 


eer 
or 
MnCl, + Ca(OH). = Mn(OH). + CaCly 
The manganese hydroxide, if exposed to the air, takes up 
OH 
oxygen and forms hydrated manganese dioxide, O=Mag 
(or MnO,.H2O). The oxidation is hastened, practically, by 
passing air and steam into the mixture. The addition of the 
oxygen to the manganese causes the hydrogen of the hydroxyl 
groups to become acid in character (see p. 206), and the com- 


104. A TEXTBOOK OF CHEMISTRY 


pound reacts with more of the calcium hydroxide to form cal- 
clum manganite, CaMnQ;: 


O—H OH O 
O=Mn€ a Cat hs O=Mn€ Ca 42H,0 
O—H OH Pee) 
Calcium 
Manganite 


The calcium manganite is insoluble and may be easily sepa- 
rated from the solution of calcium chloride. On treatment with 
hydrochloric acid it acts in the same way as a mixture of man- 
ganese dioxide and lime would. 


CaMnO; + 6 HCl = CaCl, + MnCl, + Cl. + 3 H,O 


The solution of manganese chloride may, of course, be treated 
with milk of lime, air and steam and the cycle repeated indefi- 
nitely. This process was used for many years in manufacturing 
chlorine and bleaching powder on a large scale, but it has now 
been replaced almost entirely by the electrolytic processes, 
which are simpler and more direct. 

Properties of Chlorine. Chlorine is a greenish yellow gas 
about two and one half times (44, see p. 95) as heavy as air. It 
has a characteristic odor and, even when diluted with a large 
volume of air, attacks the nose and lungs strongly, producing the 
effect of a severe cold. The best antidote is to breathe, at once, 
the vapors of strong alcohol. A larger quantity of the gas acts 
as a violent poison and may produce fatal effects. 

Water dissolves about twice its volume of the gas and for 
laboratory experiments it is usually collected by displacement 
of air, in upright jars. It is less soluble in a concentrated solu- 
tion of salt. 

Chlorine may be condensed to a liquid by cold or by pressure. 
The liquid boils at — 33.6° under atmospheric pressure and 
freezes at — 102°. The vapor pressure at 0° is 3.66 atmospheres. 

Chlorine is a very active element and forms compounds with 
all of the elements except fluorine and those of the argon family. 
With many of the elements it will combine directly and rapidly 


CHLORINE | 105 


at ordinary temperatures and with others it combincs at a much 
lower temperature than does oxygen. 

Chlorine and hydrogen combine too slowly for the rate to be 
measured, if the mixture of the gases is kept in the dark, but the 
mixture will explode at a much lower temperature than mix- 
tures of oxygen and hydrogen. If the mixture is exposed to 
diffused daylight, the elements combine slowly at ordinary 
temperatures, and if exposed to bright sunlight or to the light 
from burning magnesium, the combination is so rapid as to pro- 
duce a violent explosion. The effect is produced by light and 
not by heat, and the rays of light at the violet end of the spec- 
trum are especially effective, just as the same light rays are most 
effective in producing changes in photographic plates. Evi- 
dently the light vibrations set up or increase some kind of vibra- 
tion within the molecules or atoms of chlorine, which makes 
these more active, but we can form only a very vague idea of 
the mechanism of the action. 

Chlorine containing a minute amount of moisture will attack 
almost all of the metals, even at ordinary temperatures, and 
more rapidly on gentle warming. A strip of copper, if warmed 
and held in the gas, takes fire and burns to cuprous chloride, 
CuezCle, which melts and runs from the end of the piece. Dutch 
metal, or false gold leaf, burns with a flash. It contains copper 
and zine and gives cuprous chloride, CugCls, and zine chloride, 
ZnClo. Powdered antimony burns with brilliant flashes and 
gives antimony pentachloride, SbCl;. Phosphorus takes fire in 
chlorine, burning to the liquid phosphorus trichloride, PCls, if 
the phosphorus is in excess, or to the solid phosphorus penta- 
chloride, PCl;, if the chlorine is in excess. 

If chlorine gas is carefully dried with phosphorus pentoxide, 
it will not act on copper, iron or other metals. This may be 
shown by passing the dry chlorine into a flask containing dry 
Dutch metal. The metal will remain perfectly bright, but the 
introduction of the slightest trace of moisture will cause the im- 
mediate combination of the chlorine with the leaf. The reason 
for this catalytic effect of the water is not understood. Because 


106 A TEXTBOOK OF CHEMISTRY 


of this property, dry, liquid chlorine is kept and sold in strong 
steel cylinders. 

If a piece of tissue paper, which has been dipped in warm tur- 
pentine, CioHi¢, is thrust into a jar of chlorine, the turpentine 
will usually take fire and burn with a very smoky, red flame, 
giving hydrochloric acid and carbon : 


C10Hig - 8 Cl. = 10 Ce -f- 16 HCl 


Chlorine and Water. Bleaching. A solution of chlorine in 
water has the same greenish yellow color which is characteristic 
of the gas. Such a solution apparently contains most of the 
chlorineassuch, but it has been shown (Jakowin, Z. physik. Chem. 
29, 613) that a small amount of the chlorine reacts with the 
water : 

Cl—Cl + H—O—H = HCl + Cl—O—H 
Hypochlorous 
Acid 

The reaction is reversible, with the equilibrium far toward 
the side giving chlorine and water, and very little hypochlorous 
acid is present in the solution. If the solution is exposed to 
light, however, the hypochlorous acid decomposes in two ways, 
giving either chloric and hydrochloric acids or oxygen and hy- 
drochlorie acid : 


2 HClO + HClO = 2 HCl + HCIO; 


Chloric 
Acid 
2 HCIO = 2 HCl + O, 


It is very noticeable that the compounds of chlorine with oxy- 
gen become more stable as the amount of oxygen in them in- 
creases. As the two reactions progress, the color of the chlorine 
gradually disappears. 

If dry litmus paper or a dry piece of colored calico is placed 
in dry chlorine, the color is affected very slowly, if at all, but if 
the paper or cloth is moistened, the color will be bleached very 
quickly. The effect of the water may be in part similar to the 
action of moisture in causing chlorine to combine with metals, 


PHASES Lae yy 


but the chlorine also reacts with the water giving hypochlorous 
acid, HCIO, and this oxidizes and destroys the coloring matter. 

* Chlorine Hydrate. Phases. If chlorine is passed into water 
which is cooled to 0°, a crystalline compound, chlorine hydrate, 
Cl..8 HO, separates. If this is warmed under atmospheric 
pressure, it decomposes at 9.6°; but if the pressure of the chlo- 
rine is increased, it may exist at higher temperatures, or if the 
pressure is lessened, it will decompose at a lower temperature. 
We have, in this case, a system of three phases, solid, liquid and 
gas, which can exist over a range of several degrees of tempera- 
ture. This is true of any other system containing two com- 
ponents. Such a system with two phases may still have two 
degrees of freedom (see p. 77) — freedom to change in tempera- 
ture and freedom to change in pressure. The addition of a 
second component, chlorine, increases the number of degrees 
of freedom for a given number of phases. If the system is cooled 
to — 0.24, ice will separate from the solution of chlorine, as well 
as chlorine hydrate. When this occurs, there will be four phases 
present, liquid, ice, chlorine hydrate and vapor or gas. The pres- 
sure will also be fixed at 244 mm. No change in either tempera- 
ture or pressure can occur without the disappearance of one of 
the phases. With two components and four phases, there is no 
freedom. A further study of cases in which there are two or 
more components leads to the conclusion that the number of 
phases and number of degrees of freedom together are equal to 
the number of components increased by two, or: 


P+F=C+2 
P = number of phases 
F = degrees of freedom 
C = number of components 


This is the celebrated ‘‘ Phase rule,’ which was discovered by 
Willard Gibbs of Yale University. It applies equally well to 
the formation and decomposition of compounds, as of chlorine 
hydrate above, and to changes of state, as from ice to water and 
vapor. It is applicable only when the changes in state are 


108 A TEXTBOOK OF CHEMISTRY 


reversible and is important only when the equilibrium between 
the different phases is reached within a measurable time. 

Faraday first prepared liquid chlorine 
by warming chlorine hydrate in a bent 
tube of the form shown in Fig. 388. 
By immersing the closed, empty end, 
B, in a freezing mixture while the 
chlorine hydrate in the end, A, was 

Fig. 38 warmed, the gas liberated by the de- 

composition of the hydrate exerted 

enough pressure to cause a part to liquefy in the cooled 
end. 

The Heat of Combination of Chlorine and of Oxygen with 
Other Elements. The following are the heats of combination 
of several elements with equivalent amounts of oxygen and 
chlorine: 





H.+ O = H,0 (vapor)+ 58,000 calories 
H,+ Cl. = 2 HCl + 44,000 calories 


2Na+ O = Na,.O + 100,000 calories 
2Na+ Cl. = 2 NaCl + 195,000 calories 
Zun+- O =2Zn0O -+ 85,300 calories 
Zn+ Cl. = ZnCl, + 97,200 calories 
Cul (One ean + 37,200 calories 
Cui+, Cla = CuCl + 51,600 calories 
P,+50 = P.O; + 370,000 calories 
P, + 5 Cl, = 2 PCI + 210,000 calories 


The heat of combination with chlorine is sometimes greater, 
sometimes less, than the heat of combination with oxygen. In 
general, the heat of combination of chlorine seems to be greater 
than that of oxygen in combining with metals and less than that 
of oxygen in combining with nonmetals. 

Equilibrium in Chemical Reactions. If a mixture of four 
volumes of hydrochloric acid with one volume of oxygen is passed 
slowly through a tube containing cuprous chloride at 345°, 


EQUILIBRIUM 109 


four fifths of the hydrochloric acid will be oxidized, giving chlo- 
rine and water in accordance with the reversible reaction : 


If a mixture of equal volumes of chlorine and steam is passed 
slowly through the tube at the same temperature, one fifth of 
the chlorine will be converted into hydrochloric acid. In other 
words a mixture containing 


4 volumes or 4 molecules of HCl 
1 volume or 1 molecule of Os» 
8 volumes or 8 molecules of Cl, 
8 volumes or 8 molecules of H,O 


will be in equilibrium and will not change its composition when 
heated for a long time at 345°. We do not suppose that chemical 
action ceases, but rather that, in a given time, just as many atoms 


Z i Vols. HCl 
20 Vols. HCl Pepper eee en 2 1 Vol. O, 
5 Vols. Op ey REELED SS BT | 8 Vols. Cl, 
(8 Vols. H,O 





Fig. 39 


of chlorine unite with hydrogen to form hydrochloric acid as 
there are atoms of chlorine separated from molecules of hydro- 
chloric acid. In this way the total number of molecules of each 


4 Vols. HCl 
10 Vols. Cl, | — i Lae ei er ed rel Pe ee 1 Vol. Og 
10 Vols. H,0 | Sad irae LO eS ORR Ana ca | 8 Vols. Cle 


8 Vols. H,O 





110 A TEXTBOOK OF CHEMISTRY 


of the four substances will remain unchanged after equilibrium 
is reached, but any given atom may frequently change its state 
of combination. In such a case we may think of two opposing 
forces, one of which drives the reaction to the right and the other 
drives it to the left, and that these forces are in equilibrium. 
The force driving the reaction toward the right, when we start 
with hydrochloric acid and oxygen, must be much stronger than 
the force driving the reaction toward the left, when we start with 
chlorine and steam. This is probably due, in part, to the greater 
affinity of oxygen for hydrogen, as indicated by the heat of com- 
bination given in the last paragraph, but it is also connected with 
the change in volume which occurs in the reaction and with other 
factors which are less clearly understood. 

The reaction proceeds toward the right with the evolution of 
heat : 


1s Cs a ©) 


= H,O-+ 58,000 calories 
H, + Cl. = 2 HCl + 44,000 calories 
Hence 2HCI+0 = H.O+ Cl, + 14,000 calories 


since the sum of the reactions: 


H, + Cl, = 2 HCl 
and 2HCi+0 HO + Cl, 


| 


must give thesame amount of heat as the reaction H, + O = H.O, 
because the chlorine is in the same condition at the end as at 
the beginning. 

Whenever a reversible reaction proceeds with evolution of 
heat, a higher temperature always shifts the equilibrium in the 
direction to cause a smaller evolution of heat. In other words, 
the application of heat always helps the side of a reversible reac- 
tion in which heat is absorbed and retards that side of a reaction 
in which heat is given out. | 

In accordance with this we find that the mixture in equilibrium 
at 384° contains : 


EQUILIBRIUM 111 


4 volumes or 4 molecules of HCl 
1 volume or 1 molecule of O» 
6 volumes or 6 molecules of Cl, 
6 volumes or 6 molecules of H:,O 


This means that while four fifths of the hydrochloric acid can 
be oxidized to chlorine by the Deacon process at 345°, only three 
fourths of it can be oxidized at 384°. Itis this very unfavor- 
able effect of an increase in the temperature which makes it 
necessary to use a catalyzer and work at as low a temperature 
as possible. 

* Principle of van’t Hoff-Le Chatelier.— As has been stated 
above, an increase in temperature displaces any equilibrium in 
the direction in which heat is absorbed and an increase in 
pressure displaces any equilibrium in the direction in which the 
volume decreases. These are special cases of the principle of 
van’t Hoff-Le Chatelier, which is that every force applied to a 
system which is in equilibrium causes a change which tends to 
resist the force that is applied. Thus, if pressure is applied to 
ice, a small amount of the ice will melt, but in melting it will 
absorb heat, the temperature will fall and this will tend to 
stop the melting. Or, if dry air is blown over the surface of 
water, it will cause the water to evaporate; but as it evap- 
orates, the temperature will fall and the lower vapor pressure 
will tend to stop the evaporation. 

In accordance with the law, if hydrochloric acid and oxygen 
are brought together at 345°, heat will be evolved as they react, 
and this will tend to stop the reaction. On the other hand, if 
steam and chlorine are brought together at 345°, heat will be 
absorbed as they react, and this will tend to increase the re- 
action. 

It is necessary to distinguish between the effect of an in- 
crease in temperature to increase the speed of a reaction, 
which seems to be universal, and the tendency to cause a 
reversal of those reactions in which heat is evolved. The speed 
of the combination of oxygen and hydrogen increases rapidly 
with the temperature and becomes explosive at a very mod- 


112 A TEXTBOOK OF CHEMISTRY 


erate heat. The reverse reaction by which water dissociates 
into oxygen and hydrogen increases, however, with the tem- 
perature in accordance with the theorem of van’t Hoff-Le 
Chatelier. 

If we could understand fully the mechanism of all physical 
and chemical processes, it seems likely that we should find that 
this principle has its foundation in Newton’s law that for 
every action there is an equal and opposite reaction. 

Effect of Water on Chlorides. Ionization. When hydro- 
chloric acid or such chlorides of the metals as sodium chloride, 
NaCl, zine chloride, ZnClz, or copper chloride, CuCle, dissolve 
in water, several different lines of evidence indicate that these 
compounds separate more or less completely into chloride ions, 
‘Cl, bearing a negative charge of electricity and hydrogen, 
Ht, or metallic ions, Nat, Znt*, or Cu*t, bearing a positive 
charge, or, if the atom is bivalent or trivalent, positive charges 
of electricity. The evidence for this view is, in part, as fol- 
lows: 

1. It is found that if substances which are not electrolytes are 
dissolved in water, the freezing point is lowered in direct propor- 
tion to the amount dissolved and in inverse proportion to the 
molecular weight of the solute. An aqueous solution of alcohol 
(C2H,O, molecular weight, 46) or of sugar (CizH22O.1, molec- 
ular weight, 342) is almost as poor a conductor as pure water. 
A solution containing 46 milligrams of alcohol dissolved in 10 ce. 
of water will freeze at — 0.184°. A solution of sugar containing 
342 milligrams in 10 ce. of water will freeze at — 0.188.° One of 
these solutions must contain the same number of molecules as 
the other, and it is evident that the lowering of the freezing point 
is proportional to the number of molecules of the solute in a given 
volume of the solvent. But if we dissolve 58.5 milligrams of 
salt (NaCl, molecular weight, 58.5) in 10 cc. of water, the solution 
will freeze at — 0.349°. According to the law just stated, the 
freezing point indicates nearly twice as many molecules as there 
should be. The simplest explanation of this fact is that the 
sodium chloride separates into sodium (Na*) and chloride (CI”) 


IONIZATION 113 


ions in the solution and that these, so far as this law is con- 
cerned, act as independent molecules. 

In a similar manner, a solution containing 36.5 milligrams of 
hydrochloric acid (HCl, molecular weight, 36.5) in 10 cc. of water 
freezes at — 0.355°, indicating that it is largely separated into 
hydrogen (H*) and chloride (CI-) ions. 

2. If charged strips of metal, as, for instance, pieces of platinum 
connected with the poles of an electric battery, are dipped in a 
solution of hydrochloric acid, the chlorine atoms are attracted 
by the positive electrode (anode) and the hydrogen atoms are 
attracted by the negative electrode (cathode) and there is a 
motion of the ions throughout the solution. This is called the 
migration of the ions. The rate of migration varies for different 
ions and with the same current it is quite different, the hydrogen 
ions moving nearly five times as fast as the chloride ions. 
This effect can be shown by passing an electrical current 
from a silver anode to a platinum cathode through hydro- 
chloric acid in the U-tube (Fig. 41). A silver anode is used 
because it will combine quantitatively with the chlorine which 
is liberated, and remove it from the solution. 


The solution on this side The solution on this side 
contains : Cathode Anode contains: 
At first, cpr” At first, 





do mg. HC! = 

2 milligram atoms of H 

2 milligram atoms of Cl 

At end, 

66.5 mg. HCl = 

1.82 milligram atoms 
of H 

1.82 milligram atoms of 
Cl 

1 milligram atom of H 
is liberated 

Amount of hydrogen 

transferred across the 

linen D = 0.82 me. 

atoms. 





fa meoHcl = 

2 milligram atoms of H 

2 milligram atoms of Cl 

At end, 

43.mge..HCl = 

1.18 milligram atoms of 
H 

1.18 milligram atoms of 
Cl 

1 milligram atom of Cl 
combines with anode 

Amount of chlorine 

transferred across the 

line CD=0.18 mg. 

atoms. 


114 A TEXTBOOK OF CHEMISTRY 


At the beginning of the experiment the concentration of the 
hydrochloric acid is the same in both arms of the tube; but after 
decomposing a part of the hydrochloric acid by passing the cur- 
rent for some time, it will be found that while the concentration 
of the acid has decreased on both sides, the amount of acid on 
the cathode side has become much greater than that on the anode 
side. Since the number of hydrogen atoms liberated at the 
cathode must be exactly the same as the number of chlorine 
atoms which combine with the silver anode, the greater amount 
of acid on the cathode side must be due to the fact that the 
hydrogen ions migrate faster toward the cathode than the 
chloride ions migrate toward the anode. This will be clear from 
an examination of the figure and the accompanying statement 


a + 














about the composition of the solution at the beginning and end 
of the experiment. 

3. If a solution of sodium iodide is subjected to a powerful 
centrifugal force, the heavier iodide ions may be separated to a 
slight extent from the sodium ions (Tolman, J. Am. Chem. 
Soc. 33, 121). Similar experiments were tried with hydriodic 


IONIZATION 115 


acid, lithium iodide and potassium iodide. By whirling solu- 
tions of these substances in the apparatus shown diagrammati- 
cally in Fig. 42 the heavy iodide ions were thrown toward the 
outside, giving a negative electrical charge to the solution at 
that end of the tube and a positive charge at the inner end. 

All substances which are electrolytes are supposed to separate 
more or less into ions in aqueous solutions. Thus sodium nitrate, 
NaNO3, separates into sodium ions (Na*) and nitrate ions 
(NO;_); sulfuric acid, H,SO., may separate into two hydrogen 
ions (H+, H*), and the sulfate ion (SO, _), or it may separate, 
in part, only into a single hydrogen ion (Ht) and the acid sul- 
fate ion (HSO,-). When solutions containing electrolytes are 
mixed, the reactions which occur are, in most cases, a simple 
exchange of ions, and the groups of atoms which form the ions 
remain unbroken. Thus in the reaction: 


AgNO; + HCl = AgCl + HNOs; 


the nitrate ion (NO3_) passes from one compound to the other 
without any change. 

Such reactions are always reversible, and the equilibrium is 
frequently displaced to one side or the other because one of the 
compounds is volatile or difficultly soluble. Thus it has been 
pointed out that in the reaction : 

NaCl + H.SO, = HCl + NaHSO, 
if concentrated sulfuric acid is added to salt the equilibrium will 
be far to the right because the hydrochloric acid is a gas and 
escapes ; while if a concentrated solution of hydrochloric acid 
is added to a concentrated solution of acid sodium sulfate, the 
equilibrium may be carried to the left, because sodium chloride 
precipitates. 

Effect of Water on Chlorides. Hydrolysis. When water is 
brought in contact with a chloride of a nonmetallic element, the 
effect is very different. The chloride reacts with the water as 


though the water were composed of two parts, hydrogen, H, and 
hydroxyl, OH. 


116 A TEXTBOOK OF CHEMISTRY 


vo HOH vies: Pee 
aN + HOH =3HCl+ Pies or Cr wag 
(ol HOH OH H 
Phosphorous Acid. 
Cl OH 


Cl OH O 
Pcl +5 HOH = 5HCI+ POOH = POOH + H,O. 
Val \OH OH 
Cl OH OH 
Phosphoric Acid 


This sort of double decomposition with water is called hy- 
drolysis. ‘The hydroxyl compounds which are formed are acids. 
In most reactions between these acids and other compounds, 
the oxygen is held by the nonmetallic element, while the hy- 
drogen may be easily replaced by metals. Thus with sodium 
hydroxide we have the reaction: 


H;PO, + 3 NaOH = Na;PO, + 3 HOH 


The division between the metallic and nonmetallic elements 
in the conduct of the chlorides is not a sharp one. While sodium 
or potassium chlorides are only ionized in solution and on 
evaporation of the water the ions recombine without any loss 
of hydrochloric acid, and phosphorus pentachloride is completely 
decomposed by water and on evaporation of the solution the hy- 
drochloric acid. will escape entirely, leaving phosphoric acid, 
there are many other chlorides, like ferric chloride and aluminium 
chloride, which partly ionize and partly hydrolyze in solution. 


EXERCISES 


1. Ifa mixture of salt and manganese dioxide is treated with sulfuric 
acid, the products will be manganous sulfate, sodium sulfate, water 
and chlorine. Write the equation. 


2. How much hydrochloric acid will be required to give 10 liters of 
chlorine by the Deacon process, assuming that 80 per cent is oxidized to 
chlorine? How much hydrochloric acid will be required by the second 
stage of the Weldon process ? 


CHLORINE 117 


3. What per cent of the chlorine in hydrochloric acid is liberated 
when the acid acts on manganese dioxide? What per cent when it acts 
on potassium permanganate? What per cent when it acts on calcium 
manganite ? 

4, What is the weight of chlorine absorbed by one liter of water ? 


5. Assuming that air contains 21 per cent of oxygen (by volume), 
how many volumes of hydrochloric acid should be mixed with 100 vol- 
umes of air for the Deacon process? What per cent of free chlorine will 
the resulting mixture contain, after the reaction, if there is an oxidiza- 
tion of 80 per cent? 


CHAPTER VIII 


HYDROCHLORIC ACID. OXIDES AND OXYACIDS OF 
CHLORINE 


Hydrochloric Acid. ‘The explosive combustion of hydrogen 
and chlorine under the influence of light has been mentioned. 
Hydrogen may be burned in a jar 
of chlorine or chlorine may be 
burned in hydrogen, Figs. 43 and 
44. In each case hydrochloric 
acid is formed. It may be pre- 
pared more easily by pouring a 
mixture of 9 parts (by weight) of 
concentrated sulfuric acid with 2 
parts of water on common salt, 

Fig. 43 NaCl. Concentrated acid might 

be used, but the mixture with 
salt froths badly, while the slightly diluted acid does not froth: 


NaCl + HSO, @ NaHSO: + HCl 


Sodium Acid Sodium 
Chloride Sulfate 





Fig. 44 


The compound NaHSO, is called acid sodium sulfate because 
it still contains an acid hydrogen atom which can be replaced 
by a metal. Thus if more salt is added and the mixture is 
warmed, the reaction: 


NaHSO, + NaCl @ NaSO. + HCl 


will occur. 

These reactions are reversible and would be very far from com- 
plete in either direction if all of the substances remained mixed 
together. But as soon as the salt and sulfuric acid are mixed, 


hydrochloric acid begins to escape as a gas. When this occurs, 
118 


HYDROCHLORIC ACID 119 


the acid which has gone can no longer have any effect in driving 
the reaction in the opposite direction, and a new quantity of the 
sulfuric acid will act on the salt. In this way, if the mixture is 
warmed, the reaction may finally be made practically complete. 
The equilibrium is displaced in the direction toward the forma- 
tion of the product which is continually removed from the 
mixture. 

That the reaction is reversible may be easily shown by adding 
a concentrated solution of hydrochloric acid to a strong solu- 
tion of acid sodium sulfate. A copious precipitate of sodium 
chloride will be formed : 


HCl + NaHSO, = NaCl + H:SO, 


In this case, as the sodium chloride precipitates it can no longer 
act on the substances remaining in solution, and the equilibrium 
is displaced toward the formation of the compound which is 
precipitated. Not many years ago it was quite common to say 
that sulfuric acid is stronger than hydrochloric acid and so expels 
hydrochloric acid from its salts. The experiments described 
show that either acid may expel the other, and that the direction 
in which the reaction goes depends on the volatility or insolu- 
bility of the compounds formed and on the relative amounts of 
the reacting substances as well as upon the relative affinities 
of the chlorine and of the sulfate radical for the metal, and that 
the first three factors are frequently more important than the 
last. 

Properties of Hydrochloric Acid. Hydrochloric acid is a 
colorless gas which may be condensed to a liquid by cold and 
pressure. The liquid boils at — 83.7° and freezes at — 110°. 
What is the weight of 22.4 liters of the gas? What is the density 
as compared with air? | 

Water at 0° will absorb 503 volumes of the gas. If the solu- 
tion is boiled, more hydrochloric acid than water escapes at first, 
and the temperature gradually rises till a boiling point of 110° 
is reached. After that, the portion which distils over and that 
which remains behind will have the same composition, contain- 


120 A TEXTBOOK OF CHEMISTRY 


ing 20.2 per cent of the acid. If an acid which contains less 
than 20.2 per cent is boiled, the boiling point will be below 110° 
and the portion distilling over will contain less acid than that 
which remains. If the distillation is continued, the temperature 
will gradually rise to 110°, and after that a mixture of constant 
composition (20.2 per cent) will distill as before. 

The ratio between the volumes of hydrochloric acid and of 
the hydrogen which it contains may be demonstrated roughly 
by filling a dry tube with the gas, pouring in a few cubic centi- 
meters of liquid sodium amalgam, inserting a rubber stopper 
quickly and shaking vigorously. On opening the tube with the 
mouth below the surface of water in a beaker the water will rise 
and fill the tube one half full. Does 
the experiment demonstrate that hy- 
drochlorie acid is composed of equal 
volumes of hydrogen and chlorine? 
What would be the result if a similar 
experiment could be tried with steam 
and all of the hydrogen of the steam 
were replaced by the metal ? 

The composition of hydrochloric acid 
by volume may be demonstrated by 
the electrolysis of a strong solution of 
the acid, using carbon electrodes. The 
volumes of hydrogen and of chlorine 
liberated at the two electrodes will be 
nearly equal, if a moderately strong cur- 
rent is used and the current is continued 
till the solution around the anode is 
saturated with chlorine. A suitable ap- 

Fig. 45 paratus is shown in Fig.45. See Brown- 
lee, J. Am. Chem. Soc. 29, 237. 

For most laboratory purposes to which hydrochloric acid is 
applied the solution in water is used. The most important 
chemical properties are: 

1. Reaction with Metals. With many metals the hydrogen 





HYDROCHLORIC ACID 121 


is displaced by the metal and chlorides are formed, which dis- 
solve in the water. Thus, sodium, magnesium, zinc, iron, 
aluminium and tin give sodium chloride, NaCl, magnesium chlo- 
ride, MgCl, zine chloride, ZnCly, ferrous chloride, FeCle, alumin- 
ium chloride, AICl3, and stannous chloride, SnCle What are 
the reactions for the formation of these chlorides? It is worthy 
of notice that metals like iron and tin, which form two chlorides, 
give the lower chloride when the metals are dissolved in a solu- 
tion of hydrochloric acid. 

2. Reaction with Hydroxides of Metals. Hydrochloric acid 
reacts with hydroxides of the metals, forming chlorides and 
water : 


HCl + NaOH = NaCl + HOH 
2HCl + Fe(OH), = FeCl, + 2 HOH 
Ferrous Ferrous 
Hydroxide Chloride 
3 HCl + Fe(OH); = FeCl; + 3 HOH 
Ferric Ferric 
Hydroxide Chloride 
4HCl + Sn(OH)s = SnCl, + 4 HOH 
Stannic Stannic 
Hydroxide Chloride 


In these reactions the separation of the metallic hydroxide is 
between the metal and hydroxyl, just as the separation of the 
acid is between the hydrogen and chlorine. Compounds which 
react in this manner are called bases, the presence of a hydroxyl 
group, OH, which separates easily, being characteristic of a base, 
as the presence of hydrogen which separates easily is character- 
istic of an acid. Since hydrogen and hydroxyl have a strong 
affinity for each other and separate only to a trifling extent in 
solutions or in pure water, bases and acids newtralize each other 
by the union of the hydrogen of the acid with the hydroxy] of the 
base. The compound formed by the union of the metal with the 
chlorine or with the acid radical is called a salt. In each case, 
for the formation of a normal salt there must be as many hy- 
droxyl groups in the base as there are hydrogen atoms in the 
acid. 


122 A TEXTBOOK OF CHEMISTRY 


3. Reaction with Oxides of Metals. Some oxides of metals 
also react with hydrochloric acid to form salts and water: 


ZnO + 2 HCl = ZnCl. + H2O 


4. Reaction with Oxidizing Agents. With oxidizing agents 
hydrochloric acid is oxidized to water and chlorine, the oxidizing 
agent being at the same time reduced. In such reactions 
chlorides which contain two atoms of chlorine are in the same 
degree of oxidation as those which contain one atom of oxygen, 
since two atoms of chlorine replace one atom of oxygen in com- 
bination with hydrogen. Thus manganous oxide, MnO, is in 
the same state of oxidation as manganous chloride, MnCl, 
and ferric chloride, FeCls, corresponds in oxidation to either 
ferric oxide, Fe.O3, or ferric hydroxide, Fe(OH)3, while man- 
ganese dioxide, MnOz, is in a higher state of oxidation than man- 
ganese chloride, MnCle. 

The reactions between hydrochloric acid and manganese diox- 
ide, MnOs, potassium permanganate, KMnQsg, and calcium man- 
ganite, CaMnOs, have been given. Similar reactions take place 
with lead dioxide, PbO, and red lead, Pb30.4, which are re- 
duced to lead chloride, PbCl,; also with potassium dichromate, 
K.Cr.0;, the chromium being reduced to chromic chloride, 
CrCl;, while the potassium, which is univalent, does not change 
its state of oxidation. These equations should be written by the 
student as an aid to an understanding of reactions of this type 
and also to give practice in writing equations correctly by devel- 
oping them from a knowledge of the compounds formed instead 
of as a matter of memory. 

Indicators. A number of organic compounds are known 
which have one color in an acid solution, that is, in a solution 
containing hydrogen ions, Ht, and another color in an alkaline 
solution, that is, in a solution containing hydroxide ions, OH-. 
More strictly speaking, such compounds are, in reality, each of 
them, two different compounds so related that hydrogen ions 
will change the first into the second, and hydroxide ions will 


OXIDES AND OXYACIDS OF CHLORINE 123 


change the second into the first. Thus litmus is a red compound 
in an acid solution, and a blue compound in an alkaline solution. 
Some of the common indicators are : 








C C 
ae ah ee Nie eee: 
Litmus Red Blue 
Phenolphthalein Colorless Red 
Methyl orange Rose red Yellow 
Methyl red Red Yellow 


Congo red Blue Red 








Oxides and Oxygen Acids of Chlorine. Nomenclature 
Chlorine forms three oxides and four acids containing oxygen : 


‘Chlorine monoxide or | 


Be emernin res anhydride j 


{ ° 
C10: ~ Chlorine dioxide ~ | F010; chloe seid. 
3 


Cl,0O; — Perchloric anhydride — HCIQ, perchloric acid 


— HClO hypochlorous acid, 


The names of these acids should be learned carefully as an 
illustration of the principles used in naming acids. For chlorous 
and chloric acids the endings correspond to those which are used 
for oxides and chlorides (p. 29). The prefix hypo- means under, 
and the prefix per- means above or beyond. The endings and 
prefixes refer to the relative amounts of oxygen for different 
acids of the same element. The relations for other elements are 
not always so simple. Thus the acids of sulphur are: 


Sulfurous acid, H,SO; 
Sulfuric acid, H»SO, 
Persulfuric acid, H2,S20g [HSOu]e 
The persulfuric acid is in a higher state of oxidation than 


sulfuric acid because it contains less hydrogen, not because it 
contains more oxygen in proportion to the sulfur. 


124 A TEXTBOOK OF CHEMISTRY 


The salts of the acids are named by changing the -ous of the 
acid to -zte for the salt, and the -7c of the acid to -ate for the salt. 


Hypochlorous acid, HCIO, gives potassium hypochlorite, KCIlO 
Chlorous acid, HC1O., gives potassium chlorite, KCI1O, 
Chloric acid, HCI1Os, gives potassium chlorate, KCIO3 
Perchloric acid, HClO,, gives potassium perchlorate, KC1O4 


Hypochlorous Acid. Hypochlorites. When chlorine is dis- 
solved in water, it has been pointed out that a small amount of 
hypochlorous acid is formed by the reversible reaction : 


Cl, + HOH =@ HCl + HCIO 


The equilibrium in this reaction is very far toward the left, 
but if a base is added to thesolution, the two acids will be neutral- 
ized : 

HCl + KOH = KC] + H,0O 
HCIO + KOH = KCIO + H,O 


The neutralization of the acids causes a displacement of the 
equilibrium toward the right side of the first equation and the 
reaction goes on to completion. By adding the three equations 
together and eliminating water, the result can be expressed in the 
single equation : 


2 KOH + Cl, = KCl + KOCI + H,O 


Ii slaked lime (calcium hydroxide, Ca(OH)2) is used, a mixture 
of calcium chloride, CaCle, and calcium hypochlorite, Ca(OCl)o, 


Cl 

or a calcium chloride-hypochlorite, Ca ,is formed. This 
OCl 

is called bleaching powder. 


Ja 
Ca(OH). + Ch = Nae HO 


Calcium 
Chloride-hypochlorite 


HYPOCHLORITES 125 


The hypochlorites give up their oxygen readily to other sub- 
stances and so are powerful oxidizing agents. The action is 
much more vigorous in a faintly acid than in an akaline solution, 
however, because hypochlorous acid, HCIO, gives up its oxygen 
much more easily than a hypochlorite does. For this reason 
bleaching powder is applied to the bleaching of cotton or linen 
cloth by dipping the cloth first in a solution of the bleaching 
powder and then in very dilute acid, which liberates the hypo- 
chlorous acid. 

Not only may a hypochlorite be used to oxidize other sub- 
stances, but if a neutral or faintly acid solution of a hypochlorite 
is boiled, one portion is oxidized to a chlorate while another por- 
tion is reduced to a chloride: 


KCIO + 2KCIO = KCIO; + 2 KCl 


Potassium 
Chlorate 
If a small amount of a cobalt salt, as cobalt nitrate, Co(NQO3)o, 
is added to a solution of a hypochlorite, the oxygen of the hypo- 
chlorite is liberated in the free state. The cobalt is oxidized to 
cobalt dioxide, CoO2, which then acts as a catalyzer, as man- 
ganese dioxide acts on potassium chlorate : 


J 
2 Cac + CoO, as OP CaCl, -F O, aa CoOz 
OCI 


Hypochlorous acid is a very weak acid. While in the reaction 
of ionization : 
HCl = Ht + Cl- 
which occurs when hydrochloric acid is dissolved in water, the 
equilibrium is far to the right in moderately dilute solutions, for 
hypochlorous acid the corresponding reaction : 


Heloe Hs .0 107 
has the equilibrium very far to the left. In other words hydro- 


chloric acid gives a large proportion of hydrogen ions in dilute 
solutions, while hypochlorous acid gives only a very small pro- 


126 A TEXTBOOK OF CHEMISTRY 


portion of such ions. This fact may be used to obtain a solu- 
tion of hypochlorous acid. If hydrochloric acid is added to a 
solution of a hypochlorite, hypochlorous acid will be formed in 
accordance with the reaction : 


K* + ClO- + H+ + Clr = HClO + Kt + Cl 


From such a solution chlorine monoxide and water pass over on 
distillation. In the distillate the chlorine monoxide and water 
combine to form hypochlorous acid. An excess of hydrochloric 
acid must be avoided, however, as this would cause the reaction : 


H* + Cl + HClO @ H,0 + Cl, 


to occur, in which the equilibrium is far to the right. 

Hypochlorous acid can be obtained only in dilute solutions. 
Concentrated solutions decompose in accordance with the re- 
actions already given: 


2 HClO = 2 HCl + O2 
HClO + 2 HClO = HClO; + 2 HCl 
HCl + HClO = H.O + Cle 


* Hypochlorous Anhydride or Chlorine Monoxide. Chlorine 
monoxide, Cl.O, is formed when chlorine is passed through a 
tube containing cold, dry mercuric oxide, the mercury being con- 
verted into an oxychloride: 


2Hg0 +2Ch = HgO.HeClh, + Cl.0 


The oxide of mercury used must be obtained by precipitation, 
and washed and dried at 300° — 400°, as the crystalline oxide 
does not react readily enough. Chlorine monoxide may be con- 
densed to a liquid which boils at about 5°. Either the liquid 
or the gas explodes violently on slight provocation, giving chlo- 
rine and oxygen: 


= Cl,0 = 2.Cle ++ O. 


In this case the affinity between atoms of the same kind is 
greater, apparently, than that between chlorine and oxygen in 
chlorine monoxide. Curiously enough, when chlorine combines 


CHLORATES 127 


with a larger amount of oxygen, the compound is much more 
_ stable. 

* Chlorous Acid and Chlorites. When sodium peroxide is 
added to a solution of chlorine peroxide, ClO2, sodium chlorite 
is formed : 


2 ClO, + Na,O. ay NaClO, ++ Oz 


The chlorites are bleaching agents, similar to the hypochlo- 
rites, and are even more unstable. Free chlorous acid has not 
been prepared, even in solution. 

Chloric Acid and Chlorates. When a faintly acid solution 
of a hypochlorite is warmed, one portion oxidizes another to a 
chlorate : 


KCIO + 2 KCIO = KCIO; + 2 KCl 


Practically, if chlorine is passed into a solution of potassium 
hydroxide or into milk of lime, Ca(OH)s, till there is a slight 
excess and the solution becomes warm from the heat evolved by 
the reaction, a solution of potassium chloride and potassium 
chlorate, or of calcium chloride and calcium chlorate will be ob- 
tained. The student should write the equations and notice 
what portion of the potassium is converted into potassium chlo- 
rate. Why is it more economical to prepare calcium chlorate 
first and then obtain potassium chlorate by adding potassium 
chloride to the solution? What must be the relative solubilities 
of calcium chlorate and potassium chlorate for such a method 
to be successful ? 

Chlorie acid is stable only in solution and cannot be separated 
as a pure compound. 

* Chlorine Dioxide. When concentrated sulfuric acid is 
added to a chlorate the chloric acid liberated decomposes at once 
into perchloric acid, chlorine peroxide and water : 


KCIO; + H,SO4 = HClO; + KHSO, 
3 HClO; = HClO, + 2 ClO, + H2O 


Chlorine dioxide is a heavy, yellow gas, which is easily 
soluble in water. It seems to be even more unstable than chlo- 


128 A TEXTBOOK OF CHEMISTRY 


rine monoxide and explodes violently if warmed or brought into 
contact with organic matter. These properties may be illus- 
trated by mixing some sugar and potassium chlorate and adding 
a drop of concentrated sulfuric acid. The chlorine dioxide 
will react with the sugar and ignite the mixture. 

The solution of chlorine peroxide in water gives with a base 
a mixture of chlorite and chlorate: 


2 ClO, + 2 KOH = KCIO, + KCIO; + HO 


Chlorine dioxide may, therefore, be considered as an anhy- 
dride of both chloric and chlorous acids. 

Perchlorates and Perchloric Acid. When potassium chlorate 
is heated to its melting point, it partly decomposes into potassium 
chloride and oxygen, but a part is oxidized to potassium per- 
chlorate, while another part is reduced to potassium chloride: 


3 KCIO3 + KCIO3; = 3 KCIO, + KCl 


Potassium 
Perchlorate 


This illustrates, again, the fact that the compounds of chlorine 
with oxygen become more and more stable as more oxygen is 
taken up. This is true of the acids as well as of the salts. Per- 
chloric acid is the only oxyacid of chlorine which can be ob- 
tained as a pure compound, free from water. Perchloric an- 
hydride, Cl,OQ7, is also the most stable of the oxides of chlorine. 

It has been pointed out that concentrated sulfuric acid expels 
hydrochloric acid from salt chiefly because hydrochloric acid 
is a gas and escapes from the mixture; also that concentrated 
hydrochloric acid will precipitate salt from a solution of acid 
sodium sulfate, NaHSO,, leaving sulfuric acid in solution, chiefly 
because the salt is nearly insoluble in concentrated hydrochloric 
acid. Both principles may be used to prepare perchloric acid. 
If concentrated hydrochloric acid is poured over some sodium 
perchlorate (30 cc. for 20 grams of perchlorate), the reversible 
reaction : 


NaClO, + HCl = NaCl + HClO, 


PERCHLORIC ACID 129 


will proceed till 95 per cent of the sodium separates as sodium 
chloride. This may be removed by filtering on an asbestos filter 
and washing the salt with a small amount of concentrated hy- 
drochloric acid. The filtrate’ contains a little salt with hy- 
drochloric and perchloric acids. The highest boiling point of 
an aqueous solution of hydrochloric acid is 110°, while the boil- 
ing point of the hydrated perchloric acid is 203°. On heating 
the mixture, therefore, the hydrochloric acid escapes and finally 
the reversible reaction : 


NaCl + HClO,s = NaClO, + HCl 


is carried to completion, leaving only perchloric acid containing 
a small amount of sodium perchlorate. A pure hydrated per- 
chloric acid, containing about 28 per cent of water, may be ob- 
tained from the mixture by distilling under diminished pressure. 
This hydrated acid boils with some decomposition at 203°, under 
atmospheric pressure. Anhydrous perchloric acid may be ob- 
tained by distilling a mixture of potassium perchlorate and con- 
centrated sulfuric acid under diminished pressure. The anhy- 
drous acid is far less stable than the hydrated acid. This is 
possibly because the hydrated acid has the structure: 


O-—H 

// O—H 
H—-O-Cl Wo 
O 


* Perchloric Anhydride, Cl.O;, may be prepared by adding 
perchloric acid to phosphoric anhydride cooled to 10° and dis- 
tilling after some time: 


2 HClO, + P.O; = ClO, + 2 HPO; 
Metaphosphoric 
Acid 


Perchloric anhydride is a colorless, oily liquid which boils at 
82°. 
1 The portion of a solution which has passed through a filter. 


130 A TEXTBOOK OF CHEMISTRY 


Structure of the Oxyacids of Chlorine. If chlorine is univa- 
lent in the oxyacids of chlorine, the structure of these acids would 
be represented by the formulas : 


H—O—Cl Hypochlorous acid 
H—O—O—C]l Chlorous acid 
H—O—O—O—C] Chloric acid 
H—O—O—O—O—CI Perchlorie acid 
There is very little evidence that oxygen atoms can unite in 
this manner to form chains that are stable, and the instability 
of hydrogen peroxide, H—O—O—H, makes these formulas 
seem very improbable. There is also a good deal of evidence 
to show that chlorine and other related elements may have a 
valence as high as seven in some of their compounds. The fol- 
lowing structures are, therefore, considered much more probable : 


? 
nol H—O—Cl=0 IR boogie 4 ZO 


fo 
H—0—Cig EO 
No 

These compounds illustrate very clearly the effect of oxygen 
in giving an acid character to hydrogen compounds. Hypo- 
chlorous acid is a very weak acid and separates almost as easily 
into chlorine and hydroxyl, H—O—, as it does into hydrogen 
and —O—C]l. Perchloric acid on the contrary, is a strong, 
stable acid, especially in the hydrated form. For a possible 
explanation see p. 206. 

* The Atomic Weight of Chlorine. As oxygen is the basis for 
atomic weights, it would be most natural to determine the atomic 
weight of chlorine by determining the composition of one of the 
oxides of chlorine. But these oxides are so unstable that they 
cannot be prepared in a condition of sufficient purity for such 
a purpose. The composition of hydrochloric acid has been de- 
termined accurately, however, in two ways; first, by combining 
a weighed amount of hydrogen with chlorine, which was weighed 
in the liquid form; second, by passing a weighed amount of 


ATOMIC WEIGHT OF CHLORINE 131 


hydrogen over potassium chloroplatinate, K2PtCls, which was 
reduced to potassium chloride and metallic platinum. The 
loss in weight gave the weight of the chlorine, while the hydro- 
chloric acid formed was also collected and weighed. The aver- 
age for the ratio between hydrogen and chlorine by the two 
methods is, H: Cl = 1:35.189. As the atomic weight of hy- 
drogen is 1.0078 (p. 72), the atomic weight of chlorine is 


39.189 X 1.0078 = 35.463 


The atomic weight has also been very accurately determined 
by an entirely different method. By dissolving lithium chloride, 
LiCl, in a solution of perchloric acid and evaporating the water, 
it was converted into lithium perchlorate, LiClO,. The ratio of 
the increase in weight to the weight of the lithium chloride was : 


AO iiGle=) P50068 1936453421393 


This gives the molecular weight of LiCl as 42.393. 

Next, the amount of silver required to combine with the chlo- 
rine of the lithium chloride was determined by dissolving a 
weighed amount of silver in nitric acid and adding the solution 
of silver nitrate, AgNO;, to the solution of lithium chloride. 
This gave the ratio: 


LiCl: Ag = 0.89299: 1 = 42.393 : 107.871 


This gives the atomic weight of silver, Ag = 107.871. The 
amount of silver chloride which could be obtained from a given 
weight of lithium chloride was also determined. ‘This gave: 


LiCl: AgCl = 0.295786: 1 = 42.293 : 143.325 
Subtracting the atomic weight of silver from the molecular 
weight of silver chloride we have: 


143.325 — 107.871 = 35.454 


which is the atomic weight of chlorine. It will be seen that this 
value agrees closely with that given by the other methods, and 
it does not seem likely that the value 35.46, which is given in 
the atomic weight tables, can be far wrong. 


CHAPTER IX 


CLASSIFICATION OF THE ELEMENTS. THE PERIODIC 
SYSTEM 


Tue three elements which have been studied, oxygen, hydrogen 
and chlorine, differ very greatly from each other, but as we pass 
on to other elements, it will be found that several of these have 
very marked resemblances to chlorine, while others have proper- 
ties which recall those of oxygen, though the resemblances are 
not so close. The elements fall into a number of more or less 
well-defined families or groups, and a knowledge of these groups is 
of great assistance in acquiring a knowledge of the elements and 
their compounds. The most satisfactory classification is the one 
known as the Periodic System, which is based on the atomic 
weights and is found in the accompanying tables. The elements 
are arranged in the order of their atomic weights, with a few 
exceptions, which will be referred to below. Hydrogen and 
helium form a preliminary period and are omitted in the first 
table. ‘The valence of hydrogen would put it in Group I, while 
its amphoteric (p. 206) character in HOH would relate it to 
aluminium or silicon. Beginning with lithium, the first seven 
elements after helium, Li, 6.94; Be, 9; B, 11; C, 12; N, 14; 
O, 16; F, 19, pass from hthium, which is strongly metallic, 
to fluorine, which is very strongly nonmetallic. The first five 
form oxygen compounds as follows: LizO, BeO, B2O3, COs, 
N.O;. The last four form compounds with hydrogen, CHg, 
NHs3, OHa, FH. 

In the second row we find that sodium resembles lithium, 
magnesium resembles beryllium, and so on to chlorine, which 
resembles fluorine. Compounds with oxygen and hydrogen 


are: 
132 


CLASSIFICATION OF THE ELEMENTS 133 


Na.O, MgO, Al.Os3, S10,, P.O}, SO3, Cl.0,7 
on — SiH. PH3, SH, ClIH 


In the third row the oxides are: K,O0, CaO, Sce2O3, TiO, V2Os, 
CrO3, Mn.O,, but the last four elements do not, as in the first two 
rows, form compounds with hydrogen. At the end of the row are 
three elements, iron, cobalt and nickel, which resemble manganese 
in some of their properties. Similar groups of three elements are 
found after the fifth and seventh rows of elements. Beginning 
with the third row, the elements of alternate rows resemble each 
other much more closely than those of the successive rows, and 
each pair of rows taken together is spoken of as a long period 
to distinguish these from the short periods of the first two rows. 

It will be noticed that the highest valence of the elements 
toward oxygen increases from left to right, from one to seven: 


Ua Z Z 

“No Beeo NG cS ‘ ae va ENG 
Mea) BO a Ut en 68 No. CO 
O No Yo 

O 


The valence toward hydrogen, however, decreases from the 
center to the right from four to one: 


H 
| H H 


! te A 
H—C—H, NH, O , F-H. 
cae 


If a line is drawn in the table between beryllium and boron 
downward to the right and between tellurium and tungsten (W), 
all of the elements below and to the left of the line will be found 
to be metallic, while in the third, fifth and seventh rows the 
elements to the right of the line are also metallic. The remain- 
ing elements are nonmetallic. These facts and also the short 


1 The true formula is PsO po. 


134 


A TEXTBOOK OF CHEMISTRY 


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136 A TEXTBOOK OF CHEMISTRY 


and long periods are better shown in the form of the table on 
p. 185. The second table also shows that there is only a single 
period from cesium to radon. ‘The same relation appears in the 
curve for atomic volumes on p. 137. 

There is also a periodic relation between the atomic weights 
of the elements and their atomic volumes. The atomic volume 
may be defined as the volume occupied by one gram-atom of the 
element. Thus, the specific gravity of potassium is 0.862 and the 


0.862 





atomic volume is = 45.4. The specific gravity of silver 


107.88 


10.492 
odic relation between atomic weights and atomic volumes is 
clearly shown in Fig. 46. In general, elements with low atomic 
volumes have high melting points. This has been useful in 
selecting tungsten and tantalum for electric light filaments. 

From the time of the discovery of the Periodic Table by 
Mendeleef, it was an almost tantalizing suggestion that the 
chemical elements are all composed of the same ultimate mate- 
rials condensed in different ways. This suggestion has now 
been so far confirmed by a series of brilliant and revolutionary 
discoveries that no chemist or physicist doubts its truth. 

The most important of these discoveries have been: the 
discovery of electrons, or atoms of negative electricity, eighteen 
hundred times lighter than hydrogen atoms, by J. J. Thomson 
and Kaufmann (p. 181); the disintegration of atoms, by the 
Curies and Rutherford (p. 472); the demonstration of the 
existence of isotopes, atoms having different weights but iden- 
tical chemical properties, by J. J. Thomson, Aston, T. W. Rich- 
ards, Harkins and others; the theory of atomic structure as 
consisting of positive nuclei surrounded by rotating electrons 
~ and connected with the spectra of hydrogen, helium and other 
elements, developed by Bohr; the changes in spectra caused 
by stripping a part of the outer electrons from atoms, discovered 
by Millikan; and the actual transmutations of elements dis- 
covered by Rutherford and others. } . 





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137 


CLASSIFICATION OF THE ELEMENTS 


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138 A TEXTBOOK OF CHEMISTRY 


These discoveries have led chemists and physicists to believe 
that an isolated hydrogen atom consists of a proton, which may 
be called an atom of positive electricity, with an electron, or 
an atom of negative electricity rotating about it. A molecule 
of hydrogen consists of two protons and two electrons but the 
theories about the condition of the electrons in relation to the 
protons are not satisfactory. 

According to these new hypotheses, all other atoms contain 
a nucleus composed of protons associated with a smaller number 
of electrons than there are protons and having, for that reason, 
a positive charge. Arranged about this nucleus are successive 
layers of electrons, probably in rapid rotation. The number of 
electrons is such that in any electrically neutral atom or mole- 
cule the total numbers of the protons and electrons are equal. 
Thus a helium atom is supposed to contain a nucleus of two 
electrons and four protons, and two other electrons rotate 
about the nucleus. Chlorine consists of two isotopes: one 
having an atomic weight of 35, the other, 37. The one of 
atomic weight 35, when isolated and neutral is supposed to have 
a nucleus of 35 protons and 18 electrons. About this the first 
layer of electrons has 2, the second, 8, and the third, 7. A 
chloride ion is supposed to have 8 electrons in the outer layer, 
giving the ion a negative charge because the ion has 36 electrons 
and only 35 protons. In a molecule of chlorine two electrons 
are supposed to be shared by the two atoms, giving a total of 
70 protons, and 70 electrons. The second isotope of chlorine, 
with an atomic weight of 37, is supposed to have 37 protons 
and 20 electrons in the nucleus. Otherwise it is like the first. 

When developed in detail, this theory gives a very simple 
explanation of the periodic system and of the other new dis- 
coverles mentioned above. 

Atomic Numbers. The atomic number of an element may 
be defined as either the number of electrons outside of the 
nucleus of the neutral atom, or the number of protons in excess 
of electrons within the nucleus. 


CHAPTER X 
THE HALOGEN FAMILY 


General Properties of the Halogens. The four elements of 
the halogen family are: ! 


Fluorine, F, 19 
Chlorine, Gl 35.5 
Bromine, Br, H je80 
Todine, Ti 127 


The elements of the halogen group are the most strongly non- 
metallic of all the elements. They are also called negative be- 
cause in the electrolysis of their compounds they are attracted 
toward the anode or positive electrode. In contrast with these 
and other nonmetallic elements, metals are called positive, the 
most strongly positive or metallic elements being those of the 
alkali group, to which sodium and potassium belong. The name 
halogen means “ salt-former,” and is given to these elements be- 
cause they combine directly with metals to form salts, sodium 
chloride or common salt, NaCl, being the most important ex- 
ample. The most common salts containing other nonmetallic 
elements are those which also contain oxygen, as sodium sulfate, 
NazSOxz, or potassium nitrate, KNQO3. 

Compounds of the Halogens with Hydrogen and Oxygen. 
The compounds of the halogens with hydrogen and with hydro- 
gen and oxygen are acids and have the following formulas : 











H.F. or HE HCl HBr HI 
HCIO HBrO HIO 
—- HCI1O, Saar man 
= se HClO; HBrO; HIO; 
—— HClO, HIO, 


1 Tn this and other similar tables approximate values are given 
for the atomic weights in order that the student may learn the rela- 


139 


140 A TEXTBOOK OF CHEMISTRY 


The elements of the group are univalent in combining with 
hydrogen or with positive elements and frequently, also, in com- 
bining with nonmetallic elements. In combining with oxygen 
or with oxygen and hydrogen the valence seems to vary from | 
one in hypochlorous acid, H—O—CIl, to seven in perchloric 


O 

acid, H—0—CI€O, the odd numbers of valences being most 
O 

common. 

Fluorine is the most strongly nonmetallic or negative element 
of the group, or, indeed, of all of the elements. It will displace 
any other element of the group from combination with hydrogen 
or a metal. In a similar manner chlorine will displace bromine 
or iodine, and bromine will displace iodine. This is prob- 
ably due to the same properties which cause hydrofluoric 
acid to be the most stable and hydriodic acid to be the least 
stable of the compounds of these elements with hydrogen. 

The resemblances between chlorine, bromine and iodine are — 
much closer than the resemblances between these elements and 
fluorine. For this reason bromine and iodine are considered 
first. In studying these elements the properties of chlorine 
and its compounds should be constantly recalled and the re- 
semblances emphasized. 

Bromine, Br, 79.92. Occurrence, Preparation. In most cases 
where large quantities of chlorides are found in nature smaller 
amounts of bromides are found associated with them. In this 
way bromides are found especially in sea water and in the brines 
from which salt is obtained by evaporation and crystallization. 
Some of the American brines in Michigan are rich in bromine, 
and the bromine is obtained from these by subjecting them to 
electrolysis till all of the bromine is liberated, with a small 
amount of chlorine. As bromine boils at 59° and has a molec- 
ular weight of 160, on boiling the liquid the bromine will pass 
off with a comparatively small amount of water. (What con- 


tions among the atomic weights more easily. Accurate values are 
given on p. 10. 


BROMINE 141 


nection has the last fact with the molecular weight of bromine ”) 
The chlorine may be removed by mixing the impure bromine 
with a solution of potassium bromide and distilling : 


2 KBr + Cl, ay KCl +. Bro 


Bromine may also be prepared by warming a mixture of potas- 
sium bromide, manganese dioxide and sulfuric acid. The prod- 
ucts are potassium sulfate, manganese sulfate, bromine and 
water. What is the equation for the reaction? 

Properties. Bromine is a heavy, very dark colored liquid, 
which gives off reddish brown vapors at ordinary temperatures. 
It has a strong, disagreeable 
odor, the name having been 
given to it for this reason, 
from Bpopos, a stench. It 
is also an irritant poison. 
As with chlorine, the best 
antidote is to breathe the 
vapor of strong alcohol. 
If the liquid touches the 
skin, it produces a severe 
wound, which it is very Fig. 47 
difficult to heal. 

Although bromine vapor is much heavier than air (how many 
times heavier ?), if a little of the liquid is placed in the bottom of 
a tall cylinder, the vapor will diffuse rapidly upward through the 
air in the cylinder. How can this be explained by the kinetic 
theory ? 

Bromine combines directly with both metals and nonmetals, 
forming compounds which are, in almost all cases, very similar 
to the corresponding chlorides both in formulas and in properties. 
_ At 228° the volume of bromine vapor which would fill 22.4 
liters at 0° and 760 mm.! weighs about 160 grams, but at 1570° 
the gram molecular volume weighs only a little over 100 grams. 





1 Supposing that the vapor could be cooled to this temperature 
at a pressure of 760 mm. without its condensing to a liquid. 


142 A TEXTBOOK OF CHEMISTRY 


This indicates that at high temperatures bromine is largely 
dissociated into molecules which contain only a single atom. 
At lower temperatures the formula of bromine is evidently Bre. 

Bromine melts at — 7° and boils at 59°. It has a specific 
gravity of 3.1883 at 0° or of 2.9483 at 59°. It forms a hydrate 
which probably has the composition Bre + 8 H:O, though the 
analyses of the compound do not agree very well with the for- 
mula. Potassium and sodium bromides are used in medicine as 
sedatives, the latter by preference, because the bromide ion seems 
to be the constituent which produces the desired effect, while the 
potassium ion is much more irritant than the sodium ion when 
taken in moderate quantities. Silver bromide is used in pho- 
tography, especially in the preparation of “dry plates.” Many 
compounds of bromine are used in the manufacture of coal- 
tar dyes. Bromine has also been used to a limited extent as a 
disinfectant. 

Hydrobromic Acid. From the method used in preparing hy- 
drochloric acid we should expect to get hydrobromic acid by the 
action of sulfuric acid on sodium bromide or potassium bromide : 


KBr + H.SO, 2 KHSO, + HBr 


This reaction takes place when the substances are mixed, but 
the hydrobromic acid gas which escapes will be colored brown, 
indicating the presence of free bromine. Sulfur dioxide, SOx, 
is also found in the vapor: 


H.SO, + Be HBr Bs H.SO; ot Bre -+ H.O 
Sulfurous 
Acid 
The sulfurous acid is unstable and decomposes into sulfur 
dioxide and water: 


H.SO;3 wa H.O ++ SO2 
This is evidently because, owing to the comparatively weak 


affinity between bromine and hydrogen, hydrobromic acid acts 
as a reducing agent toward sulfuric acid. 


BROMINE 143 


To obtain hydrobromic acid free from bromine a mixture of 
hydrogen and bromine vapor may be passed through a tube con- 
taining a red-hot spiral of platinum wire. 

Another method is to drop bromine into a mixture of red phos- 
phorus and water and pass the hydrobromic acid gas through a 
tube containing red phosphorus and glass wool moistened with 
a strong solution of hydrobromic acid. ‘The method depends on 
the hydrolysis of phosphorus tribromide by water (p. 115). 


Z P + 3 Bre ae, PBr3; 
PBr3; + 3 HOH = P(OH); +3 HBr 
Phosphorous 
Acid 


Hydrobromic acid is a colorless gas, which fumes strongly in 
the air owing to its condensation with the moisture of the air 
to form a concen- 
trated solution, which 
has a much lower 
vapor pressure than 
that of water. Water 
dissolves the acid even 
more readily than it 
dissolves hydrochloric 
acid. The constant 
boiling mixture of hy- 
drobromic acid and 
water boils at 125° 
and contains 47.7 per 
cent of hydrobromic 
acid; while the cor- 
responding mixture of hydrochloric acid and water boils at 110° 
and contains only 20.24 per cent of hydrochloric acid. The 
density of the hydrobromic acid solution is also considerably 
greater for a given per cent of acid. 

Sodium Hypobromite, NaBrO, is obtained by dissolving bro- 
mine in a cold solution of sodium hydroxide, or, better, by draw- 
ing the vapor of bromine through the solution with a current 





144 A TEXTBOOK OF CHEMISTRY 


of air. (See hypochlorites.) If the solution is warmed, the 
hypobromite is changed to the bromate, NaBrQOs;: 


DNaBrO Go NaBrO > Nabe iNaproe 


Iodine, I, 126.93. Occurrence, Preparation. While a minute 
quantity of iodine is found in sea water and in almost all brines, 
the amount is too small for the practical preparation of the ele- 
ment. Many seaweeds, however, absorb a small amount of 
iodine from the sea water. The ash from these weeds is called 
kelp and contains a small amount of iodides. From these the 
iodine may be liberated by chlorine or by sulfuric acid and 
manganese dioxide. (If sodium iodide is the compound pres- 
ent, what will be the equation for the required action 7) 

Iodine is also found as sodium iodate, NaIO;3, in the crude 
sodium nitrate from Chile and Peru. The crude nitrate 
(‘‘ caliche’’) contains about 0.2 per cent of this compound and 
most of the iodine of commerce comes from this source. 

* Todine is found in the thyroid gland, and its presence seems 
to be physiologically important. The diseases of goiter and 
cretinism are, apparently, connected with a deficiency of iodine. 

Properties of Iodine. Iodine is obtained in the form of 
black, crystalline scales which melt at 114.2°. The liquid boils 
at 184.3°, but gives off a beautiful violet vapor at much lower 
temperatures. The weight of the vapor indicates that the for- 
mula is Iz at temperatures not far above the boiling point, but 
even at 700° the molecules dissociate appreciably into single 
atoms, just as the bromine and chlorine molecules dissociate 
at much higher temperatures. The stability of iodides is much 
less than that of bromides or chlorides, and the stability of the 
iodine molecule is also much less than that of the bromine mole- 
cule. In general the affinity of the nonmetallic elements toward 
metallic elements decreases with increasing atomic weight. 

Iodine dissolves very slightly in pure water. It dissolves 
more easily in alcohol, giving a brown solution, called tincture ! of 


1 The name tincture is given in pharmacy to a solution in alcohol 
of some substance or of the active constituents of some plant. 


IODINE 145 


iodine, which is used in medicine. Iodine also dissolves in a so- 
lution of potassium iodide. There is evidence that in solution it 
forms an unstable compound, KIs3, but the dilution of a solution 
having this composition causes the precipitation of a part of the 
iodine. In chloroform, carbon bisulphide and other solvents 
with which it does not combine, iodine forms violet solutions. 

Iodine gives with starch emulsion, in the presence of hy- 
driodic acid or an iodide, a deep blue color, which is very charac- 
teristic and which is used as a test for free iodine or for starch. 
From iodides the iodine must be liberated by some oxidizing 
agent, best by nitrous acid for the detection of minute quantities 
of the element. Chlorine may be used, but an excess oxidizes 
the iodine to iodic acid and destroys the color. The color of 
the starch iodide is also destroyed by heat, but returns, in part, 
on cooling the solution. 

Hydriodic Acid. It has already been pointed out that the 
affinity between hydrogen and iodine is much less than that 
between hydrogen and chlorine or bromine. If slightly diluted 
sulfuric acid is poured on potassium iodide, some hydriodie acid 
is liberated, but the larger part of the acid acts upon more of the 
sulfuric acid, reducing it to sulfur dioxide or even to hydrogen 
sulfide, H2S. The student should write the equations for the 
three reactions involved in the last statement and compare 
with the somewhat similar conduct of hydrobromic acid. 

Hydriodic acid is best prepared by melting together in a 
distilling bulb 1 part of red phosphorus with 20 parts of iodine, 
forming a mixture of phosphorus triodide, PI3, and iodine.1 
When the mixture is cold, a stopper bearing a separatory funnel 
containing 4 parts of water is fitted to the neck of the bulb, and 
the side tube is connected with a small U-tube containing a very 
little water to wash the gas and retain nearly all of the iodine 


1 Method of Lothar Meyer slightly modified, Ber. 20, 3381. 
The usual direction, which gives enough phosphorus to form PIs, 
gives rise to the formation of phosphonium iodide, PHI, and this 
may stop the exit tubes or contaminate the product. For the same 
reason iodine cannot well be removed from the gas by moistened red 
phosphorus, as directed by some authors. 


146 A TEXTBOOK OF CHEMISTRY 


which passes over. If a solution of hydriodic acid in water is 
desired, the delivery tube should not dip beneath the surface of 
the water which is to absorb the gas. 

When all is ready, the water is dropped slowly on the mixture 
of phosphorus triodide and iodine. After all has been added 
the last of the hydriodic acid is driven over by warming. The 
equation is: 


PI3 + I, + 4 H.O a H3PO, + 5 HI 
Phosphoric 
Acid 
The solution of hydriodic acid in water has properties similar to 
those of the corresponding solutions of hydrochloric and hydrobro- 
mic acids. The solu- 
tion of constant boiling 
point under atmos- 
pheric pressure boils 
at 127°, has a specific 
gravity of 1.70 and 
contains 57 per cent 
of the acid. The hy- 
driodic acid is slowly 
oxidized with libera- 
tion of iodine on ex- 
posure of the solution 
to the air, and for this 
reason the aqueous 
Fig. 49 acid is almost always 
colored red or brown. 
Direct Combination of Hydrogen and Iodine. Reversible re- 
actions. Equilibrium. Mixtures of equal volumes of chlorine 
and hydrogen or of bromine vapor and hydrogen combine com- 
pletely when heated to the temperature of combination, and 
neither hydrochloric acid nor hydrobromie acid dissociates ap- 
preciably unless heated to a quite high temperature.! If a mix- 





1 Haber (Thermodynamik technischer Gasreaktionen, 1905, S. 
95) calculates the dissociation of hydrobromie acid as only 0.15 per 
cent at 727°. 


EQUILIBRIUM 147 


ture of equal volumes of hydrogen and of iodine vapor is heated, 
however, there is no temperature at which the combination to 
form hydriodic acid will be complete, even if the mixture is 
heated for an indefinitely long time. On the other hand, if 
hydriodic acid is heated, it decomposes slowly, even at quite 
low temperatures, but never completely, no matter how long it 
is heated, unless the temperature is very high indeed. If we 
start with two sealed glass tubes, one containing one part by 
weight of hydrogen with 127 parts of iodine and the other 
containing hydriodic acid, and heat both tubes at the same tem- 
perature till the composition no longer changes, it will be found 
that each tube contains, on cooling, a mixture of hydrogen, 
iodine and hydriodic acid but that the composition of the mix- 
ture in the two tubes is identical. This result is most easily 
explained by supposing that we have here a reversible reaction : 


Hele aH eee E 


and that when equilibrium is reached the reaction does not stop 
but continues in such a way that just as many molecules of hy- 
driodic acid are formed in a minute as are decomposed in the 
same time. 
The composition of the mixture which is in equilibrium 
varies with the temperature, as will ke seen from the following 
table: 


COMPOSITION OF THE EQUILIBRIUM M1IxTuRE oF HYDROGEN, 
IopINE AND Hypriopic AcID 











Warne ents Meee PROPORTION OF PROPORTION OF HyDROGEN 
Hypriopic Acip AND IopINE 
283° 0.8213 0.1787 
328° 0.8115 0.1885 
374° 0.7990 0.2010 
393° 0.7942 0.2058 
427° 0.7843 0.2157 


508° 0.7592 0.2408 





148 A TEXTBOOK OF CHEMISTRY 


It is evident from this table that when equilibrium is reached, 
the combination of hydrogen and iodine is more nearly complete 
at low temperatures than at high ones. (In accordance with the 
principle of Le Chatelier (p. 111) is the combination accompanied 
by the evolution or by the absorption of heat ?) 

Speed of Chemical Reactions. If hydriodic acid is heated at 
a given temperature, the decomposition seems to proceed more 
and more slowly as the reaction goes on, as will be seen from the 
following table: 


Rate oF DECOMPOSITION OF Hypriopic AciIp at 374°! 














TotTaL TIME OF FRACTION FRACTION DECOMPOSED 
HEATING DECOMPOSED IN 1 MINUTE 
360 minutes 0.0715 0.00020 
720 minutes 0.1267 0.00015 
1080 minutes 0.1596 0.00009 
1440 minutes 0.1715 0.00003 





A little consideration of these results leads us to the conclusion 
that the decreasing rate is due to two causes: first, because the 
amount of hydriodic acid in the mixture is constantly decreasing, 
and second, because the hydrogen and iodine which result from 
the dissociation are recombining. 

If we begin with the mixture of hydrogen and iodine, the com- 
bination appears to be rapid at first, but soon decreases in its rate 
as the amounts of hydrogen and iodine grow less and as the 
hydriodic acid which is formed begins to decompose. In order 
to get the real rate for the decomposition of pure hydriodic acid 
or for the combination of pure hydrogen and iodine, we might 
measure the rate for the first infinitesimal fraction of a minute, 
if that were possible. For the decomposition at 374° the deter- 
mination of the composition of the mixture at the end of six 
hours gives a rate for the decomposition of 0.0002, or 1/5000 


1 Bodenstein, Z. physik. Chem. 29, 295. 


SPEED OF REACTIONS 149 


part of the whole in one minute, and this is comparatively close 
to the rate of decomposition for pure hydriodic acid, since at 
the end of the first six hours, only about 1/15 of the whole has 
been decomposed. 

* Concentration and Speed of Reaction. In order to calcu- 
late the true rate more accurately, it is necessary to use the law of 
the relation between the concentration of reacting substances and the 
speed of reaction,’ which has been based on a careful study of 
many different reactions which take place slowly enough so that 
the rate can be measured. This is that the speed of any reaction 
at a constant temperature (2.e., the part of the whole which will 
react in unit time) zs equal to the product of the concentrations of 
each reacting substance multiplied by a force which is characteristic 
of the given reaction. 

The force which causes the reaction, and which is given a 
numerical value in relation to the speed of the reaction by this 
law, is spoken of in most textbooks as chemical affinity, but it is 
evidently complex, depending on the attraction between the 
atoms which unite, the attraction between atoms which separate, 
the temperature, the presence of catalytic agents and perhaps 
on many other factors. The temperature, especially, has a large 
effect, such that the speed of a reaction is usually doubled for 
an increase of 10°. 

For hydriodic acid the law may be given the following expres- 
sion : 

Let Cy, = Concentration of hydrogen, 
Cy, = Concentration of iodine, 
Cyr = Concentration of hydriodic acid, 


F, = Force driving the reaction, H, + I, == HI + HI, to 
the right, 


1 Often called, less correctly, the ‘‘Law of Mass Action.” 

2 In the case of hydriodic acid the nature of the force is wholly 
ae by the action of light and the reaction becomes unimolec- 
ular : 

HI=H4+I. 


Bodenstein, Z. physik. Chem. 22, 23. 


150 A TEXTBOOK OF CHEMISTRY 


F, = Force driving the reaction to the left, 
S; = Speed of formation of hydriodic acid, 
S2 = Speed of decomposition, all at a given temperature 
Then: 
Cu, X Cr X Fi = S81 
Cur X Cur X Fo = Se} 


By means of these formulas it is possible to calculate from the 
results obtained by heating mixtures of hydrogen and iodine, 
or by heating hydriodic acid for different lengths of time, the rate 
of combination or of decomposition for the pure substances at 
unit concentration. The methods of calculation are compli- 
cated and need not be given here. The results are as follows: 


Rate oF FoRMATION AND DECOMPOSITION OF Hypriopic AcID 




















S 
TEMPER- St FRACTION on WHOLE Si 
ATURE ela ee. ae WHOLE DECOMPOSED IN ONE So 
RMED IN ONE MINUTE Nikos 
302° 0.000353 0.00000326 1: 108 
374° 0.0140 0.000221 I hig een ess 
427° 0.172 0.0031 1D 








It is clear from this table that the equilibrium mixtures as given 
on p. 147 contain much more hydriodic acid than hydrogen 
and iodine because the combination takes place much more 
rapidly than the decomposition. It is also seen that the speed 
of each reaction increases rapidly with the temperature and that 
the speed of decomposition increases more rapidly than the speed 
of combination. From this it follows that combination is ac- 


1 This depends on the fact that the decomposition results from 
the action of two molecules of hydriodic acid on each other with the 
formation of molecules of hydrogen and iodine. Such a reaction 
is called bimolecular. The action of light seems to cause the direct 
separation of the atoms of hydrogen and iodine from each other 
and the reaction in that case is unimolecular. See footnote, p. 149. 


SPEED OF REACTIONS 151 


companied by an evolution of heat, in accordance with the 

principle of Le Chatelier (see below and also pp. 111 and 201). 
* Calculation of the Relative Speed of Two Reactions from 

the Composition of an Equilibrium Mixture. If we know the 

nature of a reversible reaction and the composition of the equilib- 

rium mixture, it is possible to calculate the relative speeds of 

the opposing reactions, at unit concentration. 

From the formulas given: 


Cy, X Cr, X Fi = 81 
Cut X Cur X Fo = S82 


From these equations, F, = S; and F2 = S2 when the concen- 
tration is 1 for each of the reacting substances. 
At 374° the composition of the equilibrium mixture is very 
near to: 
80 per cent of HI, 
10 per cent of He, and 
10 per cent of Iz, by volume. 


From this: 
Cur = 0.8 
Cu, = 0.1 
Crm Oe 


2 


Let S3be the speed of combination and S, the speed of decomposi- 
tion, at equilibrium. At equilibrium S; must equal S,. 
Then, 
Cu, x C1, x F, = S3 
0.1X 0.1 X Fi = 83 
Cur X Cu X Fe = S4 
0.8 X 0.8 X F. = S4 
Since S3 =s S4 
vin 
2 
it should, with the result obtained by the direct measurement 
of the speeds of the two reactions (p. 150). The result may 


and = 64. It will be seen that this result agrees closely, as 


152 A TEXTBOOK OF CHEMISTRY 


also be checked by the following calculation, putting Fi = Si 
and F, = S: and using the values for S; and S2 given on p. 150: 


Cy, X Cr X Si = 83 

0.1 X 0.1 X 0.0140 = S3 = 0.000140 
Cur X Cur X So = Sa 

0.8 X 0.8 X 0.00022 = S4 = 0.000140 


This means, of course, that after equilibrium is reached at 
374°, the dissociation and recombination still continue at the 
rate of about 1/7000 part of the whole 
each minute. 

Effect of Removing one of the Reacting 
Substances. Displacement of the Equi- 
librium Point. If a tube containing hy- 
driodie acid is heated in such a way that 
one end of the tube is kept cool, the iodine 
which results from the dissociation will partly condense and the 
concentration of the iodine, Cj,, will be diminished in the equation : 


Cy, X Cr, X Fi = 83 


which gives the rate of recombination. Under these conditions 
it is evident that the decomposition must go much farther than 
usual before S; = S4. In other words, the removal of one of the 
constituents of a reversible reaction always displaces the equilib- 
rium to the side on which the constituent removed appears. 
This effect has been noticed before in the reaction between 
iron and steam and in that between salt and sulfuric acid. If 
the end of the tube could be kept cold enough so that the vapor 
pressure of iodine in it would be reduced to zero, S3 would finally 
become zero, and the decomposition of the hydriodic acid would 
continue till it was complete. 

Heat of Formation of Hydriodic Acid. The heats of formation 
of the compounds of the halogens with hydrogen are as follows: 


H,+ F. = 2 X 38,000 calories 
H, + Cl, = 2 X 21,800 calories 
H, + Bre. = 2 X 8300 calories 
H,+I1I, =2xX 96 calories 





FLUORINE 153 


The heat energy liberated during the combination grows less 
with increasing atomic weight and becomes very small in the 
case of iodine. 

Bodenstein calculates (Z. physik. Chem. 29, 313) the heat of 
combination of hydrogen and iodine in ‘gaseous form as follows : 


At 510°, 2 & 2222 calories 
At 290°, 2 X 943 calories 
At 20°,2 xX 96 calories 


Fluorine, F, 19.0. Occurrence. The chlorides, bromides and 
iodides of four of the most common metals, calcium, magnesium, 
sodium and potassium, are all easily soluble in water, and these 
three halogens are found chiefly associated with these metals, and 
especially with sodium, in the ocean and in brines. Fluorine, on 
the other hand, combines with calcium to form an almost insolu- 
ble compound, calcium fluoride, CaFs, and for this reason can 
never be found in more than very small amounts in natural 
waters, which practically always contain calcium. Fluorine is 
found chiefly as calcium fluoride, CaF., in the mineral flworite. 
Cryolite, a double fluoride of aluminium and sodium, NasAlF's 
(or AlF3.8NaF), found in Greenland, and apatite, a double phos- 
_ phate and fluoride, or chloride, of calcium, Cas(PO,)3F or 
Cas(PO,)3Cl,! found in Canada and elsewhere, are important 
as sources of aluminium and phosphorus rather than as sources 
of fluorine. 

Preparation. When we consider that the heat of combination 
of hydrogen and fluorine is nearly twice that for the combination 
of hydrogen and chlorine (p. 152), and remember that chlorine 
can take a part of the hydrogen away from oxygen, we may be 
led to expect that free fluorine cannot exist in the presence of 
water. It was not till this came to be clearly understood that 
Moissan succeeded in obtaining the free element in 1886. He did 
this by electrolyzing a solution of potassium fluoride, KF, in 
anhydrous hydrofluoric acid. He used a U-tube of platinum 


1 Note the relation between this formula, the formula of phosphoric 
acid, H3PQO,, and the valences of calcium and of fluorine or chlorine. 


154 A TEXTBOOK OF CHEMISTRY 


at first, but showed later that a tube of copper is only very 
slightly attacked by the fluorine, if the temperature is kept at 
— 23°, or below, by a freezing mixture. In the electrolysis, 
fluorine goes toward the anode and is liberated there while 
potassium and hydrogen go toward the cathode, but only hydro- 
gen is liberated, because hydrogen ions are discharged at a much 
lower potential than potassium ions. 

Properties. Fluorine is a greenish yellow gas, less deeply 
colored than chlorine. The weight of a gram molecular volume 
is 38 grams, showing that the formula is Fy. Fluorine is the 
most active of the nonmetallic elements, as is to be expected 
from its unique position in the periodic system. It combines 
directly and vigorously with nearly all elements, both metals and 
nonmetals, except with oxygen. Many elements, as iodine, 
phosphorus, arsenic, carbon as charcoal or lampblack, silicon, 
potassium and sodium take fire and burn in the gas, forming 
fluorides. Fluorine will also displace nearly all other non- 
metallic elements from their compounds. If led into water, 
it gives hydrofluoric acid and oxygen, rich in ozone: 


3H, + 3F,. = O3 + 6HF 
Ozone Hydrofluoric 
Acid 


Etching Glass. Hydrofluoric Acid may be easily prepared 
by warming a fluoride with concentrated sulfuric acid : 


CaF, + H.SO4 = CaSO, + 2 HF 


Hydrofluoric acid is a gas which may be condensed to a liquid 
much more easily than the other halogen acids. The anhydrous 
liquid boils at 19.4°. The liquid mixes with water in all propor- 
tions, the concentrated solution fuming in the air in the same 
manner as concentrated solutions of the other halogen acids. 
The gaseous acid 1s very poisonous and the concentrated or anhy- 
drous acid causes painful wounds, which are very difficult to 
heal. 

The most interesting property of the acid is its action on sili- 


FLUORINE 155 


cates and especially on glass, which is a complex silicate of 
calcium and sodium or other metals. When hydrofluoric acid 
comes in contact with glass, the fluorine combines both with the 
silicon and with the metals of the glass: 


CaSiO3 + 6 HF = SiF, + CaF, + 3 HO 
Calcium Silicon 
Silicate Tetrafluoride 


The reaction may be looked on as a displacement of oxygen 
by fluorine, two atoms of fluorine displacing one atom of oxygen 
in accordance with the valences of the two elements. Silicon 
tetrafluoride, Sify, is a gas and escapes. By covering a glass 
object with beeswax, which is not affected by hydrofluoric acid, 
and exposing it to the action of the gas, after drawing lines or 
figures through the wax so as to expose part of the surface of the 
glass, it is possible to etch the exposed surface and obtain per- 
manent markings of anyform that is desired. Graduation marks 
on thermometers, burettes, eudiometers, etc., are made in this 
way. The best results are obtained by exposing the glass to 
the anhydrous gas for some hours. 

Commercial hydrofluoric acid is kept in lead bottles, which are 
only slightly attacked. The pure acid must be kept in platinum 
or in bottles made of ceresin, a mineral wax with a higher melting 
point than that of paraffin. The constant boiling solution boils 
at 120° and contains 35 per cent of the acid. 

Hydrofluoric acid, unlike the other halogen acids, forms both 
acid and neutral salts. Thus it forms with potassium, acid 
potassium fluoride, KHF%, as well as the neutral, or normal, fluo- 
ride, KF. The formation of these acid salts seems to be closely 
related to the abnormal density of the gas and indicates that the 
true formula of the acid in solution or at low temperatures is 
probably H.F; instead of HF. The weight of a gram molecular 
volume of gas varies from 51.2 grams at 26° to 20.6 grams at 88°. 
At the lower temperature the gas is evidently more complex 
than H2F:, for which the gram-molecular-volume would weigh 
40 grams. 


156 A TEXTBOOK OF CHEMISTRY 


Metallic Elements of Group VII. Manganese stands between 
chlorine and bromine in the seventh group of the Periodic System 
when the system is given its simplest form (p. 134). It re- 
sembles chlorine in the dioxide, MnOe, which corresponds to 
chlorine dioxide, CLOs, and in permanganic acid, HMnOs,, 
corresponding to perchloric acid. But in most of its properties 
manganese is metallic, and it will be considered further later 
(p. 533). 

The Periodic System indicates the possibility of three or four 
other elements in the seventh group with atomic weights greater 
than that of bromine, but no such elements have been dis- 


covered. 
EXERCISES 


1. Write the equations for sixteen reactions between the following 
acids and bases, giving normal salts : Hydrochloric acid, HCl; perchloric 
acid, HC1O,; sulfuric acid, H2SO,; phosphoric acid, H3PO,; sodium 
hydroxide, NaOH; ferrous hydroxide, Fe(OH)2; ferric hydroxide, 
Fe(OH)3; stannic hydroxide, Sn(OH),. 

2. Write the equations for the reactions between the following salts 
and sulfuric acid: sodium chloride, NaCl; calcium chloride, CaCl; 
sodium perchlorate, NaClO,; aluminium chloride, AICI. 

3. Write the equations for the reactions between hydrochoric acid 
and the following oxidizing agents. Notice the changes in valence: 


MnO, > MnCl, 

KMnQ, — KCl and MnCl, 
HClO +>HCl 

Pb30, —-> PbCl. 
K.Cr.0; > KCI and CrCl; 
KCIO; ~> KCl 


4. Write the equation for the reaction between potassium iodide, 
manganese dioxide and sulfuric acid, giving K2SO, and MnSQ,. 

5. Write the equation for the reaction between calcium bromide, 
CaBre, potassium permanganate, KMnO, and sulfuric acid, giving 
calcium sulfate and the other products to be expected. 

These reactions are introduced here to give the student facility in 
writing equations on the basis of the valence of the elements. The 
fundamental conception of valence is that each atom has the power 
of holding directly in combination a definite, small number of 
other atoms. Thus, when we write the graphical formula H—Cl, 
the thought which it is intended to convey is that a hydrogen or 


THE HALOGEN FAMILY 157 
a chlorine atom holds directly to only a single other atom Fy the 


compound, hydrochloric acid. In water, H—O—H or is 


oxygen won holds directly to two other atoms. In some sense we 
may think that an oxygen atom has two points of attachment for 


ing A Aine é 
other atoms. In ammonia, N@H or HNC , in accordance with 
S H 


the same theory, each nitrogen atom holds directly to three hydrogen 
atoms. 

In the series of oxides of the Periodic System the elements of the zero 
group do not combine with other elements at all, and these elements 
are considered to have a valence of zero. The elements of the first 
group are univalent, and one bivalent oxygen atom can hold two atoms 
of these elements as in Na.O or Na—O—Na._ The bivalent atoms of 
the second group can hold bivalent oxygen atoms, atom for atom, as 
in Mg=O. In the third group, where the elements are trivalent, if 
we consider an atom of such an element as combined with one atom of 
oxygen, one valence of the first element will remain unsatisfied, thus, 


BC. If a second atom of oxygen is added, one valence of this will 


GO 
be unattached, B\g. On adding a second atom of the trivalent 


element and a third atom of oxygen, all of the valences will be balanced 
Be> 

yO. In the next group a quadrivalent atom can balance two bi- 
BXO 


valent oxygen atoms, BG: The same principles may be easily 


extended to the compounds, N2O;, SO; and Cl,0;. 

When the oxides are those of nonmetallic elements, they will, in most 
cases, combine with water to form acids. In this case one valence of 
one oxygen atom separates from one nonmetallic atom, and the hydro- 
gen, H, of the water attaches itself to the oxygen, while the hydroxyl, 
OH, of the water attaches itself to the nonmetallic atom : 


O O 
nZO n&o . : 
og O-H O H- 
AS er a or o=s€ —> s¢ 
2 /O-H O H-O O 
N=O N<=O 


158 A’ TEXTBOOK OF CHEMISTRY 


In the reactions between acids and bases the same principles of bal- 
ancing valences are to be applied, the only difference being that the 
valences of the metal on the one hand are to be balanced against the 
valences of the acid groups on the other. Since the hydroxyl group 
—QO—H is univalent, the number of hydroxyl groups in the base gives 
the valence of the metal of the base, while the number of replaceable 
hydrogen atoms gives the valence of the acid group. Thus iron is bi- 
valent in ferrous hydroxide, Fe(OH)s, and trivalent in ferric hydroxide, 
Fe(OH);,! while the sulfate group, SOu, of sulfuric acid is bivalent and 
the phosphate group, PO, of phosphoric acid, H3PQk,, is trivalent. By 
representing the valences with lines, it is a simple matter to balance 
the valences of a metal against the valences of an acid radical and so 
determine the correct formula of a salt. Thus for ferric sulfate the 
formula must be: 


FeZ SOx 
so, 

Pet 
NSO; 


or for ferric phosphate, Fe=PQ,. A little practice of this sort will 
soon enable a student to write correct formulas, such as Fee(SO,)3 or 
FePQ,, without the use of the lines to indicate valences. If these prin- 
ciples are once understood, one needs to remember only the formula of a 
single salt of any metal with some well-known acid in order to be able 
to write the formulas of the normal salts of the metal with a hundred 
or more acids whose formulas are known. 

In reactions which involve oxidation and reduction it often happens 
that the valence of some element changes. In a reduction, oxygen 
or some other element is removed without being replaced, or hydrogen 
is added, and to do this hydrogen must usually be furnished from 
some source, and the element combined with this hydrogen is often 
liberated in the free state. 

Thus in the reaction between manganese dioxide and hydrochloric 
acid quadrivalent manganese changes to the bivalent form. The extra 
oxygen atom is balanced by hydrogen from the hydrochloric acid and the 
chlorine of the latter is liberated. The manganese dioxide is reduced, 
the hydrochloric acid is oxidized: 


OF Re cl 
Zea Cl We 
Mn€ +5 c= Ma --2H.0— Cl; 
He 


1 For the sake of simplicity the possibility of such doubled 
formulas as Fe,(OH), and Fe(OH). is not presented here. 


THE HALOGEN FAMILY 159 


In a similar way if hydrochloric acid acts on potassium permanganate, 
KMnQ,, only three chlorine atoms are taken by the potassium and 
manganese, and the hydrochloric acid which furnishes these will give 
only three of the eight hydrogen atoms necessary to balance the four 
oxygen atoms of the permanganate molecule. To balance the remainder 
of the oxygen atoms, five more hydrogen atoms will be required. This 
gives us the reaction : 


KMnOQ, + 3 HCl = KCl + MnCl, + 4H,0 + 5Cl 
+ 5 HCl 


If we wish to take account of the fact that free chlorine has the 
formula Cle, the equation must, of course, be doubled, giving : 


2 KMnO, + 16 HCl = 2 KCl + 2 MnCl, + 8 H.0 + 5 Cl,’ 


Which substance is reduced and which is oxidized in this reaction ? 


CHAPTER XI 
SULFUR, SELENIUM AND TELLURIUM 


THE nonmetallic elements of Groups VI and VII of the peri- 
odic system are: 


Ode heb Boys ese ee LG 
Sh ueteskeeie wore CUS eo ieao.o 
DCm ue od wie Bree ou 
Tell Pia 6 Donhos bE gnaw 


Sulfur, S, 32.06. Occurrence. Oxygen is found free in nature, 
partly because of its great abundance, forming, as it does, one 
half of that portion of the earth which we can examine directly, 
partly, probably, because of its unique relationship to carbon 
and the growth of plants (p. 312). Sulfur is also found free, 
partly because it is a comparatively abundant element and 
partly because it is easily liberated from hydrogen sulfide and 
other sulfides by the action of oxygen and some compounds of 
oxygen. Free sulfur is found in large quantities in Sicily and in 
Louisiana. Until about 1903 the sulfur mines of Sicily held, 
for a long time, a practical monopoly of the sulfur markets of 
the world, almost the only competition coming from the sulfur 
obtained by the Chance process (p. 457) as a by-product in the 
manufacture of sodium carbonate. The sulfur in Sicily is mixed 
with other minerals, from which it is separated by piling up the 
mixture and setting fire to the sulfur in such a way that the 
heat from burning a part of the sulfur melts the rest and the 
latter runs out and is collected. The process is, of course, a 
wasteful one —as a pound of coal would give nearly as much 
heat as four pounds of sulfur (p. 27). The crude sulfur is re- 
fined by distillation. If the vapors are condensed in cold cham- 
bers, the sulfur takes the form of flowers of sulfur, just as the 


freezing of water vapor gives snow. If the condensing room is 
160 


SULFUR 161 


above the melting point of sulfur, the liquid sulfur which collects 
on the bottom is run into molds and forms the roll brimstone of 
commerce. 

The extensive deposits of sulfur in Louisiana are below a layer 
of quicksand, and for a long time after they were discovered no 
practical method of working the deposits was known. The 
difficulty was finally solved by a process invented by Mr. Frasch 
of New York. ‘Three concentric iron pipes are sunk to the level 
of the sulfur, and hot water under pressure is forced down be- 
tween the two outer pipes, the 
pressure of the water being great 
enough so that the boiling point 
is raised above the melting point water 
of the sulfur, 114.5°. The hot 


{ air 


ws 


| being SUIfur 










: “hy Z YU be 
water melts the sulfur, which yy j hi Mist? 
rises in the second tube, the end = // Yj Lyi, 
of which is brought below the sur- y Uy YD 
face of the melted sulfur. To Z LAD), 
bring the sulfur to the surface, air U, WY l 
is forced down through the cen- 6 Ye 
tral tube, the sulfur and com- Whe 
pressed air rising together between I Py pst: 
the central and second tubes. U Li LaL Ld 


By this process the production 
of sulfur in the United States was 
increased from 3500 tons in 1900 
to 265,000 tons in 1911. The 
world’s production of sulfur in 1909 Yj sulfur Yf 
was 818,000 tons. dditittitize Y 

Sulfur is also found in nature Fig. 51 
combined with metals as metallic 
sulfides and with metals and oxygen as sulfates. The most 
important sulfides are lead sulfide, or galena, PbS, zine sul- 
fide, or sphalerite, ZnS, iron sulfide, or pyrite, FeSe, and an 
iron-copper sulfide, copper pyrites, CuFeS2. The most important 
sulfates are calcium sulfate, or gypsum, CaSO4.2 HO, and barium 






| water 






tL, 
eo 


162 A TEXTBOOK OF CHEMISTRY 


sulfate, or barite, BaSO,. Of these, only iron pyrites is used 
primarily as a source of sulfur, for the manufacture of sulfuric 
acid. The other sulfides are used primarily for the metal which 
they contain, but sulfuric acid is sometimes made from them as 
a by-product. 

Allotropic Forms of Sulfur. Sulfur may exist in three well- 
defined solid forms, in two liquid forms, which correspond 
closely to two of the solid forms, and in three gaseous forms. 
The solid forms are: 

1. Rhombic Sulfur. Light yellow crystals, most often in the 
form of rhombic pyramids (p. 194), found in nature and formed 
by crystallization from carbon disulfide, in which sulfur is easily 
soluble. The specific gravity is 2.06 and the melting point 
114.5°. This is the most dense and most stable form at ordinary 
temperatures, and the other forms change to this form more or 
less rapidly at temperatures below 96°. 

2. Monoclinic Sulfur. When melted sulfur is allowed to 
cool slowly, it crystallizes in long, transparent needles of the 
monoclinic system (p. 195). These have a specific gravity of 
1.96 and melt at 119°. This form of sulfur is stable only at tem- 
peratures between 96° and 119°. At lower temperatures it 
changes more or less quickly to the rhombic form. The outer 
form of the needles is retained, but they become opaque and then 
consist of microscopic crystals of the rhombic form. 

3. Amorphous, Insoluble Sulfur. When sulfur which is 
heated above 160° is cooled quickly with care that it dees not 
come in contact with crystals of sulfur, which would cause a 
rapid transformation to the crystalline form, it assumes a soft, 
plastic form, which hardens to a solid mass after some hours 
or days. If this hardened mass is treated with carbon disulfide, 
it will be found to be mostly insoluble and the insoluble portion 
is amorphous, 1.e. it has no crystalline structure. 

The liquid forms of sulfur are: 

1. Mobile Liquid Sulfur (S,). Between the melting point 
(114.5° or 119°) of either form of sulfur and 160° it forms a 
mohile, pale yellow liquid. 


SULFUR 163 


2. Viscous Liquid Sulfur (S,). When heated to 160°, sulfur 
suddenly becomes dark colored and so viscous that a test tube 
containing it may be inverted without its running out. If 
heated to a higher temperature, the liquid becomes gradually 
somewhat more mobile and finally boils at 444.7°.1. The boiling 
point is frequently used to fix a point on the scale of ther- 
mometers and pyrometers. 

The gaseous forms of sulfur are: 

1. Ss. When sulfur is converted into a vapor at 250°, under 
low pressure, the weight of a gram molecular volume is nearly 
256 grams, indicating that there are eight atoms in one molecule 
and that the formula is Sg. 

2. Se. Even at the boiling point (444.7°), the weight of a gram 
molecular volume of sulfur vapor is considerably less than 256 
grams and it was formerly supposed that the formula at tempera- 
tures a little higher than this was Ss. A more careful study of 
the matter has demonstrated that the formula Sg is the true one 
at low temperatures and that the heavy molecules dissociate 
as the temperature rises until, at 800°, the formula becomes So, 
the weight of a gram molecular volume at that temperature 
being 64 grams. It is still somewhat uncertain whether the 
larger molecules dissociate directly into molecules of S2 or 
whether intermediate molecules of S4 or Sg are found.? 

3. S. When sulfur vapor is heated to a very high tempera- 
ture (2000°), it dissociates still further until the gram molecular 
volume weighs only 32 grams and the formula becomes 8. We 
may suppose that at high temperatures the collisions between 
molecules become more and more violent until, at last, the 
affinity between the atoms can no longer withstand the disrup- 
tive effect of the collisions. 

Properties and Uses of Sulfur. Sulfur burns readily in air 
or oxygen, forming sulfur dioxide, SOs, with usually a small 
amount of the trioxide, SO3. The volume of the sulfur dioxide 


1 Bulletin of the Bureau of Standards, Vol. 7, pp. 3 and 129. 
2 See Prenner and Schupp, Z. physik. Chem. 68, 144 (1909), and 
Stafford, ibid. 77, 66 (1911). 


164 A TEXTBOOK OF CHEMISTRY 


is almost the same as the volume of the oxygen from which it is 
formed. (How does this follow from Avogadro’s law and the 
formulas of oxygen and sulfur dioxide?) Sulfur combines with 
most metals when heated with them, forming sulfides. The 
combination with iron to ferrous sulfide, FeS, and with copper 
to cuprous sulfide, Cus, is attended with considerable evolution 
of heat. 

Sulfur is burned to sulfur dioxide for the manufacture of sul- 
furic acid, for use in bleaching straw goods, for the “ sulfuring ”’ 
of fruit in the process of drying, to prevent darkening and the 
growth of harmful organisms. Sulfur is also used in the manu- 
facture of carbon disulfide, of gunpowder and of india rubber. 
It is used directly or in a lime-sulfur wash for application to 
vines, fruit trees, etc., to prevent the growth of fungi or other 
harmful organisms. 

Hydrogen Sulfide, H2S, is found in many natural waters, the 
so-called sulfur waters. It is formed by the decomposition of 
organic matter containing sulfur and is one cause, though by no 
means the only reason, for the disagreeable odor of decayed eggs 
and sewage. . 

Hydrogen sulfide is formed when hydrogen is passed over 
sulfur heated to its boiling point, as can be shown by passing the 
gas, subsequently, through a solution of lead nitrate, in which it 
will produce a black precipitate of lead sulfide: 


Lead Lead 
Nitrate Sulfide 


The reaction is reversible: 
2 He + So = 2 HS 


as can be shown by passing hydrogen sulfide through a hot glass 
tube, in which a ring of sulfur will be deposited beyond the point 
that is heated. 

Hydrogen sulfide is prepared in the laboratory by the action 
of hydrochloric or sulfuric acid on ferrous sulfide. 


HYDROGEN SULFIDE 165 


FeS + 2 HCl = FeCl, + HS 
Ferrous 
Chloride 
FeS + H.SO, == FeSO, a HS 
Ferrous 
Sulfate 
For the preparation of the gas on a small scale the apparatus 
used for the preparation of hydrogen is suitable. For the use 
of a laboratory the Parsons apparatus (J. Am. Chem. Soc. 25, 
233) is better, because 
the acid remains in con- 
tact with the ferrous 
sulfide till the action is 
complete. Hydrochloric 
acid is more satisfactory 
than sulfuric acid for 
such a generator, be- 
cause ferrous sulfate is 
less soluble than fer- 
rous chloride and some- 
times crystallizes in the 
tube through which 
the spent acid escapes 
(Pig’-52). 
Hydrogen sulfide is a 


colorless gas with a very | ___ ORE 













x 
% 
Ay os 






MW 
R 
A} 






jizz 
a 
I 


Se 
ele 





disagreeable odor. It 
is quite poisonous, if 
breathed in more than 
small amount. It may 
be condensed to a liquid, : 
which boils at — 62° and 
frozen to a solid, which Fig. 52 
melts at — 85°. 

Solution of Hydrogen Sulfide. Henry’s Law. One volume of 
water absorbs, or dissolves, 4.4 volumes of hydrogen sulfide at 




















166 A TEXTBOOK OF CHEMISTRY 


0°, 3.7 volumes, at 10° and 3.1 volumes at 20°. The volume of 
the gas dissolved is, between quite wide limits, independent of 
the pressure. Since the weight, or amount of the gas in a given 
volume, is proportional to the pressure, it follows that the 
amount of the gas dissolved varies directly with the pressure. 
This is known as Henry’s Law (discovered in 1803). It applies 
to partial pressures also. Thus if a gaseous mixture contains 
10 per cent by volume of hydrogen sulfide, the amount dissolved 
from such a mixture at 20° will be only 0.31 volume. One hun- 
dred cubic centimeters of water in contact with pure oxygen dis- 
solve 4.9 cc. of the gas at 0°; in contact with nitrogen 100 ce. 
dissolve 2.35 cc. of nitrogen. One hundred cubic centimeters 
of water in contact with air will contain, therefore, 4.9 X 0.21 
= 1.04 ce. of oxygen and 2.35 X 0.78 = 1.83 cc. of nitrogen. 
The law does not hold for gases which are very easily soluble 
in water, such as hydrochloric acid or ammonia. 

In accordance with Henry’s law, water containing hydrogen 
sulfide loses the gas rapidly on exposure to the air, in which the 
partial pressure of the gas is, of course, zero. In addition to 
this the oxygen absorbed by the water reacts with the hydrogen 
sulfide, liberating sulfur : 


The action is similar to the liberation of chlorine from chlorides 
by fluorine, but is far less rapid. 

Sulfides. Groups of Analytical Chemistry. When hydrogen 
sulfide is passed into a neutral or slightly acid solution containing 
salts of certain metals, such as arsenic, mercury and lead, the 
metal is precipitated as a sulfide because the sulfides of these 
metals are extremely insoluble: 


2 AsCls3 + a H.S — AsoS3 4+ 6 HCl 
HgCl., + HeS = HgS + 2 HCl 
Pb(NO3)2 + H2S = PbS + 2 HNO3 
If hydrogen sulfide is passed into an alkaline solution contain- 
ing the salts of some other metals, such as iron, zinc and manga- 
nese, which are not precipitated from acid solutions, these metals, 


STRENGTH OF ACIDS 167 


whose sulfides are also very insoluble, but more soluble than 


those of the metals of the first group, are precipitated also as 
sulfides. 


FeSO, + HeS + 2 NaOH (or NaS) = FeS + Na.SO, + 2 H,O0 
ZnSO.1 + NaoS = ZnS + NaeSO, 
MnSO, + NaS = MnS + Na.SO, 


A part, but not all, of the metals of the first group are precipi- 
tated from alkaline as well as from acid solutions. The reason 
for the exceptions need not be discussed here. (See p. 261.) 

A third class of metals form salts which are not precipitated 
from acid, neutral or alkaline solutions. 

The conduct of solutions of metals toward hydrogen sulfide, 
as just outlined, is the basis for the separation of metals into three 
fundamental groups for the purposes of analytical chemistry. 

Hydrosulfuric Acid. Strength of Acids. A solution of hydro- 
gen sulfide in water will redden blue litmus paper and will neutral- 
ize a solution of sodium hydroxide, or in other words a certain 
amount of a solution of a base must be added before the hy- 
droxide will turn the litmus blue. These are the properties of 
an acid, and hydrogen sulfide is sometimes very properly called 
hydrosulfurie acid, just as hydrogen chloride is called hydrochloric 
acid. We have seen that a solution containing a milligram 
molecule of hydrochloric acid in 10 cc. of water freezes at 
— 0.355°, while a solution of alcohol containing a milligram 
molecule in 10 ce. freezes at — 0.184, and the difference was 
explained by supposing that the hydrochloric acid separates 
largely into hydrogen (H*) and chloride (Cl-) ions. A solution 
of hydrogen sulfide which contains one milligram molecule in 
10 cc. freezes at — 0.196°. 


36.5 mg. HCl in 10 ce. of H2O freezes at — 0.355 
46 mg. CeH,O in 10 ce. of H2O freezes at — 0.184 
34 mg.He2S _ in 10 cc. of H:O freezes at — 0.196 


This indicates that such a solution of hydrogen sulfide con- 
tains comparatively few hydrogen ions. This conclusion is 


168 A TEXTBOOK OF CHEMISTRY 


confirmed by the electrical conductivity of the solution. The 
solution of hydrochloric acid referred to is a very much better 
conductor (nearly 2000 times) of electricity than the solution of 
hydrogen sulfide. 

During a long period in the history of chemistry acids were 
spoken of as strong or weak according to whether they could 
expel other acids from their salts or not. Thus sulfuric acid was 
thought to be stronger than hydrochloric or nitric acid because 
it would expel these acids from salt or saltpeter. We have seen 
that such a view can no longer be held (p. 119) and that all such 
reactions are reversible. There is another sense, however, in 
which some acids are strong while others are weak, and the basis 
for a true distinction of this kind has just been indicated. A 
strong acid is one which separates largely into hydrogen ions and 
negative ions in an aqueous solution. A weak acid is one that 
separates to only a comparatively small degree into hydrogen 
ions and negative ions in solution. In this sense hydrochloric 
acid is one of the strongest of the acids, sulfuric acid is weaker 
but still a very strong acid, hydrofluoric acid is much weaker, 
acetic acid is still weaker and hydrosulfuric acid, H2S, is very 
weak indeed. 

Acids like hydrosulfurie acid which contain two hydrogen 
atoms may ionize in either of two ways: 


H.S = Ht + HS 
or HS =” HHS ae 


In the case of weak acids the ionization probably takes place 
almost exclusively in the first form. It is worthy of notice that 
the halogen acids (hydrochloric acid, ete.), which contain only 
one hydrogen atom in the molecule, ionize very completely in 
moderately dilute solutions, while hydrogen sulfide, with its two 
hydrogen atoms, ionizes to only a slight extent. If oxygen is 
added, however, as in sulfuric acid, H.SOu,, the ionization be- 
comes large, though it does not equal that of hydrochloric acid. 

Application of the Idea of Strength of Acids to explain the 


a ee ee 


STRENGTH OF ACIDS 169 


Conduct of Sulfides.1_ Practically all of the ordinary reactions 
in aqueous solutions are reversible. A reversible reaction leads 
to a stable condition only when the reaction has reached a point 
where it proceeds just as fast in one direction as in the other. 
In the reversible reactions : | 


Pb**+ +2NO0O;-+ Ht+HS = Pb. +2 NO; -+ H* and 
SH 


pM HS se Pbs + HS? 

\sH 
the reactions will proceed toward the right as long as lead sul- 
fide (PbS) is formed and separates from the solution. The 


equilibrium finally reached must depend upon whether there 
are enough lead ions (Pb*t*), lead hydrosulfide (Pb ) ions and 
SH 
hydrosulfide ions (HS~) in a given volume of the solution to form 
more lead sulfide than can remain in solution. As lead sulfide 
is very insoluble, only a very few hydrosulfide ions can remain 
in a solution containing lead ions. 
The ionization of hydrogen sulfide : 
Hes Pid Ht + HS— 

takes place to a very limited extent even in pure water, the 
equilibrium in the reaction of ionization being very far to the 
left. If we add to a solution of hydrogen sulfide a strong acid, 
as hydrochloric acid, which gives a large number of hydrogen 
ions, the hydrosulfide ion (HS~) will meet hydrogen ions more 
frequently than before and will combine with them to form 
hydrogen sulfide. This must shift the equilibrium to the left 
and cause an increase in the unionized hydrogen sulfide and a 


1The student should read this paragraph, but it may be well to 
leave its careful study till review or a later period. (See pp. 379-386. ) 
2 This may involve the further ionization, 
Pbt en Bhs +P Ht 


but the ion HN if capable of existence at all, would immediately 
become PbS. Most authors are accustomed to write the reaction: 
Pbt++ 2NO;- + 2H+ + S-- = PbS + 2 NO3- +2 H+, but the 
form given above seems more probable. 


LEO A TEXTBOOK OF CHEMISTRY 


decrease in the number of hydrosulfide ions! (HS~). The 
presence of a moderate amount of hydrochloric acid in a solu- 
tion containing a lead salt will, therefore, so far decrease the con- 
' centration of the hydrosulfide ions that lead sulfide can no longer 
be precipitated. 

It will be seen from what has just been said that the distinction 
between the first and second classes of metals (p. 166) in qualita- 
tive analysis depends on our definition of » “ slightly acid ”’ 
solution. Such metals as lead, cadmium and zine might belong 
to the first class or the second according to the concentration 
of the hydrogen ions present. 

The addition of an alkali, as sodium hydroxide (NaOH), to a- 
solution of hydrogen sulfide has an effect opposite to the addition 
of an acid. The base gives hydroxide (OH) ions, which combine 
with the hydrogen ions to form water. This displaces the 
equilibrium for the ionization of hydrogen sulfide in the oppo- 
site direction and results in a large increase in the number of 
hydrosulfide ions (HS~). Under these conditions the sulfides 
of iron, zinc and some other metals, which are too soluble to 
form at all in acid solutions, will form and be precipitated. 

When hydrogen sulfide is passed into a solution containing 
a hydroxide of a metal of the third class, a hydrosulfide, which 
ionizes to a large extent in dilute solutions, is formed : 


Nat + OH- + Ht + HS” = Nat + HS + H.O 


If a second, equal amount of sodium hydroxide is added and 
the solution is evaporated to dryness, sodium sulfide, Na.S, 
may be obtained : 


Nat + HS” + Nat + OH” = Na.S + H;O 


1 This may be stated mathematically as follows: 

Cyt X Cus~ X Fi = Si 

Cus X Fo = Se» 
Since S; = S, at equilibrium and S, must be constant for a given 
quantity of hydrogen sulfide and F; is also constant, the product 
of Cut X Cus~ must be constant for any given concentration 
of hydrogen sulfide. Any increase in the number of hydrogen ions 
must, therefore, be accompanied by a corresponding decrease in the 
number of hydrosulfide ions. 


REDUCTION BY HYDROGEN SULFIDE 171 


If the sodium sulfide is dissolved in water, the ionization of 
water approaches so near in degree to that of hydrogen sulfide 
that the sodium sulfide is largely hydrolyzed : 


Nass -+ Ht + OH” = 2 Na* + HS~ + OH- 


The presence of hydroxide ions in such a solution is indicated 
by the alkaline reaction of the solution, as shown by litmus or 
other test papers. 

Hydrogen Sulfide as a Reducing Agent. Hydrogen sulfide 
readily gives up its hydrogen to chlorine, bromine or iodine. 
It also gives up hydrogen to a great variety of compounds, re- 
ducing them. ‘The following are typical illustrations : 


Hs+1.=2HI4+5 


This reaction furnishes an excellent method of preparing a solu- 
tion of hydriodic acid, by suspending iodine in water and passing 
hydrogen sulfide into the mixture. 


K.Cr.0, + 8 HCl + 3 H.S = 2 KCI + 2 CrCl; + 7H.O+ 38 
Potassium Chromic 
Dichromate Chloride 


In writing the equation for this reaction, notice that the formulas 
of the chlorides determine the number of molecules of hydro- 
-chioric acid required. Comparing the number of molecules of 
hydrochloric acid with the number of atoms of oxygen in the 
dichromate it is seen that after water has been formed from the 
hydrogen of the hydrochloric acid three atoms of oxygen remain. 
These will oxidize the hydrogen of three molecules of hydrogen 
sulfide. 
Feo(SOx)3 +- H.S == 2 FeSO, he H.SO, + S 
Ferric Ferrous 
Sulfate Sulfate 
Here the hydrogen of the hydrogen sulfide takes the sulfate 
radical (SO,) from the ferric sulfate, and the iron is reduced from 
the ferric to the ferrous state. 
Another method of writing such equations, which is preferred 
by some teachers, is based on the principle of positive and nega- 
tive valences. According to this principle : 


172 A TEXTBOOK OF CHEMISTRY 


1. The algebraic-sum of the valences of any compound 1s zero. 
The valence of a free element is also zero. 

2. Oxygen in compounds has a negative valence of 2. 

3. Hydrogen in compounds has a positive valence of 1. 

4. When the valence of one element changes, the valence of 
some other element or elements must change by the same amount 
in the opposite direction. 

In applying these principles to the present case, the equation 
is first written in the following form: 


K.Cr2O, -}- HCl + H.S —> 2KCl ++ 2 CrCl, a S a H,O 


On inspection it is seen that the sum of the valences of the two 
potassium and two chromium atoms on the left is + 14, while 
on the right the sum of the valences of the same four atoms is 
only + 8, a loss of 6 positive valences. To balance this the 
valence of the sulfur atom changes from — 2 in hydrogen sul- 
fide, H.S, to O in free sulfur, S. It is obvious, at once, that 
to balance the changes in the chromium we must have three 
molecules of hydrogen sulfide. To furnish the chlorine for the 
chlorides there must be 8 molecules of hydrochloric acid, HCl. 
The equation becomes, therefore : 


Similar reactions occur between hydrogen sulfide and chlorine, 
hydrogen sulfide with sulfuric acid and potassium permanganate, 
KMn0O,, or hydrogen sulfide and ferric chloride, FeCl; The 
reaction between lead sulfide, PbS, and nitric acid, HNOs, 
giving lead nitrate, Pb(NOs3)2, nitric oxide, NO, sulfur and 
water, is also closely related to these, the positive, bivalent lead 
atom taking the place of the two positive hydrogen atoms in 
hydrogen sulfide. The student is advised to write the equa- 
tions for these reactions by use of both of the methods suggested 
above. 

Sulfur Dioxide is formed when sulfur is burned in the air, 
also when iron pyrites, FeSe, is burned, the latter method of 
preparation being used largely in the manufacture of sulfuric 
acid. 


SULFUR DIOXIDE 173 


Sulfur dioxide may be prepared in the laboratory by the reduc- 
tion of concentrated sulfuric acid with copper or other sub- 
stances. If copper is used, copper sulfate is formed. The 
equation may be written as follows: 


Cu -f- H.SO, — [CuO] + SO. + H.O 
[CuO] + H2SO4 = CuSO, + HO 
Combining, Cu + 2 HeSO, = CuSO, + SO. + 2 H.O 


The first two equations are written as an aid to the writing of 
the last. The [CuO] is placed in brackets to indicate that it is 
not a final product of the reaction. It may or may not be formed 
as an intermediate product. Another method of writing would 
be to represent hydrogen [2 H] as an intermediate product. 

The most convenient laboratory method for the preparation 
of sulfur dioxide is to drop concentrated sulfuric acid into a 40 
per cent solution of acid sodium sulfite, NaHSO;: 


NaHSO; -f- H.SO, Fa NaHSO, oo H.SOz3 


Sulfurous 
Acid 


HicO, = SO, HG 


Sulfurous acid, H2SO3, decomposes very easily into sulfur 
dioxide and water, the sulfur dioxide escaping as a gas. This 
has the same effect on the equilibrium of the first reaction as if 
the sulfurous acid itself were volatile and escaped from the mix- 
ture. 

Sulfur dioxide is a colorless gas with a suffocating odor, famil- 
iar in the burning of sulfur-tipped matches. It may be con- 
densed to a liquid in a tube surrounded with a freezing mixture 
and boils at — 10°. It freezes at a very low temperature and 
melts at — 73°. 

Sulfur dioxide is used to bleach straw, wool and silk. The 
latter, especially, are injured by the action of chlorine, so that it 
cannot be used. The sulfur dioxide seems to combine with the 
coloring matter to form colorless compounds, or, in some cases, ° 
to reduce the colored compound to a colorless one. Exposure 


174 A TEXTBOOK OF CHEMISTRY 


to the air and light frequently restores the color, as in the case 
of straw hats. 

Sulfur dioxide and sulfites are powerful germicides. Its use 
as a disinfectant, however, has been almost entirely replaced by 
formaldehyde, which is even more effective and does not injure 
fabrics or metallic articles, as sulfur dioxide does. The injury to 
fabrics may be either through its bleaching effect or because 
it is slowly oxidized by the action of air and moisture to sul- 
furic acid, which is corrosive. Sulfur dioxide is still used exten- 
sively in “ sulfuring ” fruit to destroy the organisms which cause 
darkening and injury during the drying. 

Sulfurous Acid. At ordinary temperatures water dissolves 
about 50 times its volume of sulfur dioxide. The solution red- 
dens litmus and neutralizes bases, showing that the sulfur dioxide 
combines with the water and forms an acid. The sodium salt 
obtained by neutralizing the acid is sodium sulfite, Na2SOs, 
and from the formula of this and other salts it is assumed that 
the formula of the sulfurous acid in such a solution is H2SQs. 
The structure of the acid is probably 


Ne vy 
o% oH 


It is very unstable, one hydroxyl group and one hydrogen atom 
separating very easily from the molecule. The solution smells 
strongly of sulfur dioxide, and all of the gas can be expelled by 
boiling the solution. 

Sulfurous acid is a comparatively weak acid. In a solution 
containing 0.05 gram molecule (= mol) about 20 per cent of 
its hydrogen is ionized,! while in a corresponding solution of 
sulfuric acid 60 per cent of the hydrogen is ionized. Sulfurous 
acid is a powerful reducing agent. It is oxidized to sulfuric acid 


1 This is on the supposition that all of the sulfur dioxide has 
combined with water to form sulfurous acid. It is probable that 
some of the sulfur dioxide’ exists as such in the solution and that the 
ionization of the sulfurous acid really present is considerably greater 
than appears from these figures. 


_, 


SULFUR TRIOXIDE 175 


by potassium permanganate, KMnQOu,, potassium dichromate, 
K.Cr.0;, chlorine, bromine or ferric salts. 

Sulfites. Sulfurous acid forms both acid and normal salts, 
the salts of sodium being acid sodium sulfite, NaHSOs, and nor- 
mal sodium sulfite, Na,SO3. The calcium salts are CaH2(SO3)o 
and CaSO3. These salts are prepared, commercially, by burn- 
ing sulfur in air and passing the mixture of sulfur dioxide 
and nitrogen through a solution of sodium carbonate or sodium 
hydroxide for the sodium salts, or through milk of lime (Ca(OH ).) 
for the calcium salts. The acid sodium salt has been used as an 
addition to wine or cider to stop fermentation. ‘The acid cal- 
cium salt is used in the purification of wood pulp for the manu- 
facture of paper. 

Sulfur Trioxide. Some heat is evolved when sulfur dioxide 
combines with oxygen to form the trioxide, SO3, but the speed 
of the reaction between the two is too slow to be measured at 
ordinary temperatures. At temperatures where the speed of 
the reaction of combination becomes sufficiently rapid to become 
a practicable method of preparation, the dissociation of sulfur 
trioxide into sulfur dioxide and oxygen becomes very large and 
renders this method of preparation from the substances alone 
impracticable. The reversible reaction : 


2 SO. + O2 = 2 SO3 


has its point of equilibrium shifted toward the left as the tem- 
perature rises, in accordance with the principle of van’t Hoff- 
Le Chatelier (p. 111). 

As early as 1831 it was discovered that the reaction is greatly 
accelerated by the presence of platinum, but it was nearly 
70 years before the details for the application of this principle 
were so far worked out as to render the manufacture on a 
large scale possible,! so long does it often take to convert a 
scientific discovery into commercial success. The chief diffi- 
culties to be overcome were, first, that arsenic and other sub- 


1 For very interesting historical details see Knietsch, Ber. 34, 
4069 (1901). 


176 A TEXTBOOK OF CHEMISTRY 


stances in the gases obtained by roasting pyrites “ poison ” 
the platinum and render it ineffective for the catalysis, and 
second, that the platinum catalyzes the dissociation of sulfur 
trioxide as well as the combination of sulfur dioxide and oxygen, 
and the temperature for rapid combination lies very close to a 
temperature at which the dissociation is large and so the com- 
bination becomes incomplete. These difficulties have been 
overcome by a careful purification of the sulfur dioxide as it 
comes from the pyrites burners and by a careful regulation 
of the temperature as the gases pass over the “ contact mass.” 
The platinum of the “ contact mass” is disseminated in a very 
finely divided condition over asbestos or some other material 
which gives it a large surface in proportion to its weight. 

In the laboratory, on a small scale, sulfur trioxide can be read- 
ily prepared by passing dry sulfur dioxide and oxygen through a 
gently warmed tube containing platinized asbestos (Fig. 53). 





It may be obtained still more easily by warming “ fuming ” 
sulfuric acid, which is a mixture of sulfuric acid, H2SO,, pyro- 
sulfuric acid, H2S.O;, and sulfur trioxide. 

Sulfur trioxide is a clear, volatile liquid which solidifies at a 
low temperature. It melts at 14.8° and boils at 46°. In the 
presence of a trace of moisture a little sulfuric acid, H2SQu, or 
pyrosulfuric acid, H2S20,, is formed and this acts as a catalytic 
agent causing sulfur trioxide to polymerize, forming the com- 
pound S20., which crystallizes in white, asbestos-like needles. 
As it is extremely difficult to exclude moisture completely, this 
polymeric form is usually obtained instead of the true trioxide. 
On warming it gives a vapor, which consists of the true trioxide. 


SULFURIC ACID 177 


Sulfur trioxide hisses like a hot iron when thrown into water, 
owing to the heat generated when it combines with water to 
form sulfuric acid. It fumes strongly in the air, forming minute 
drops of sulfuric acid, which settle only very slowly and are not 
readily absorbed by water. Curiously enough these minute 
drops are easily absorbed by concentrated sulfuric acid and this 
is used for the purpose in the manufacture of sulfuric acid by 
the contact process. 

Sulfuric Acid. The contact process for the preparation of 
sulfur trioxide has, thus far, been used almost exclusively for the 
manufacture of a very concentrated or a “fuming”’ sulfuric acid. 

It has been pointed out that the direct combination of sulfur 
dioxide and oxygen is too slow to be commercially possible as a 
method of manufacture, and that platinum is used to catalyze, 
or hasten, the reaction. Another catalytic agent, not so sensitive 
to impurities in the gases, or to temperature changes, and which 
acts rapidly at ordinary temperatures, has been used for a long 
time in what is called the “‘ chamber process ”’ for the manufac- 
ture of sulfuric acid. In this process large chambers lined with 
sheet lead, which is only slightly attacked by dilute sulfuric acid, 
are employed. Into these chambers are introduced : 

1. Sulfur dioxide from burning sulfur or iron pyrites: 


2 FeS. + 11 O = Fe.03 + 4 SO, 
2. Nitric acid from Chile saltpeter, NaNOs, and sulfuric acid : 
NaNO; + H.SOs = NaHSO, + HNO; 


3. Air, to furnish oxygen. 

4. Water as steam or spray. 

The first reaction consists in the oxidation of the sulfur dioxide 
to sulfuric acid by the nitric acid: 


1 This equation should not be learned by rote, but should be written 
on the following considerations : 

1. When nitric acid is reduced to nitric oxide, two molecules 
give 3 atoms of available oxygen. 

2. Each atom of oxygen will oxidize one molecule of sulfur 


dioxide. 


178 A TEXTBOOK OF CHEMISTRY 


If it were necessary to stop here and the nitric oxide were lost, 
sulfuric acid would be, comparatively, an expensive substance on 
account of the limited supply and relatively high price of sodium 
nitrate. But nitric oxide combines almost instantly with the 
oxygen of the air to form nitrogen dioxide, NO: 


2 NO + O2 = 2 NO, 


Nitrogen dioxide, in turn, can oxidize a new quantity of sulfur 
dioxide : 
SO. + NO, + H.O = H.SO, + NO 

It is pretty certain that the mechanism of the reaction is more 
complicated than is indicated by these equations, but the equa- 
tions given indicate clearly the fundamental facts on which the 
action depends. These are: first, that nitric oxide combines, 
practically instantaneously, with oxygen; second that nitrogen 
dioxide can, directly or indirectly and very quickly, give its 
oxygen to the sulfur dioxide and water, converting these to 
sulfuric acid. Commercially, the whole process depends on the 
speed with which these actions occur. 

The theory of the lead chamber process which has received 
most acceptance is that of Lunge, who supposes the process to 
consist in the formation and decomposition of nitrosyl sulfuric 


acid : ONO 
ce 
no” 


a mixture of nitric eae NO, and nitrogen dioxide, NOg 
(equivalent to nitrous anhydride, N2O3), being the effective 


agent. O—NO 
2 SO. + NO; + O. + H.O = 2, s02<. 
OH 


JPNO JOH 


Nitrosyl sulfuric acid is a definite, crystalline compound, 
which is formed in the chambers when the supply of water is 


SULFURIC ACID 179 


insufficient, but it exists only as an intermediate product, if at 
all, in the normal manufacture. Nitrous anhydride, also, can 
exist only momentarily, if at all, as it decomposes at once into 
nitric oxide and nitrogen dioxide at the temperature of the 
chamber. For a further discussion of the subject see Trautz, 
Z. physik. Chem. 47, 513; Wentzki, Z. angew. Chem. 23, 1907 ; 
Raschig, zbid. 23. 2241, 24, 160; Ber. zbid. 23, 2250. 

If it were possible to lead into the chamber pure oxygen and 
sulfur dioxide, a small amount of nitric acid would convert an 
indefinitely large amount of sulfur dioxide into sulfuric acid. 
Since, however, air containing only 21 per cent of oxygen mixed 
with 79 per cent of nitrogen (and argon) must be used, there must 





Gay-Lussac 


Pyrites Leaden Chambers Tower 


Burners 


Fig. 54 


be a constant escape of nitrogen, carrying with it nitric oxide 
or nitrogen dioxide at the further end of the chamber or set of 
chambers. To recover these the gases are passed through a 
tower, known as the Gay-Lussac tower, Fig. 54, in which they 
are exposed to a large surface of concentrated sulfuric acid 
running down over broken coke or a series of earthenware plates. 
The strong acid absorbs the oxides of nitrogen, forming nitrosyl 
sulfuric acid, SO.(OH)(ONO). This nitrated acid is then forced 
by compressed air to the top of another tower, called the Glover 
tower, B. Here it is mixed with some of the more dilute 


180 A TEXTBOOK OF CHEMISTRY 


acid from the chamber and a little nitric acid to replace the un- 
avoidable loss. The mixture runs down over broken coke and 
comes in contact with sulfur dioxide coming from the pyrites 
burners, C. This causes the denitrification of the acid: 


2 SO.(OH)(ONO) + SO, + 2 H.0 = 3 H.SO, + 2 NO 


The nitric oxide is, of course, carried back into the first 
chamber. When these towers are used, only from 25 to 40 
pounds of sodium nitrate are required for the manufacture of 
a ton of sulfuric acid. Without them, two or three times as 
much is required. 

The acid from the chambers has a specific gravity of 1.53 to 
1.62, and contains 62-70 per cent of the pure acid. It is usually 
concentrated to about 79 per cent by evaporation in lead pans. 
At this point the acid begins to attack the lead more strongly 
and the concentration is completed to a specific gravity of 
1.83-1.84 and 93 to 95 per cent, in glass, platinum or iron, the 
last metal being only slightly attacked by the concentrated acid, 
although it dissolves easily in the dilute acid. If the concen- 
trated acid is distilled, an acid of constant composition contain- 
ing about 98.5 per cent of the pure acid finally passes over at 
338°. The density of the vapor proves that the process is not 
ordinary boiling, but consists in the dissociation of sulfuric acid 
to sulfur trioxide and water and that the two recombine on 


cooling : 1.801 = SO; 4+ H,0 


The specific gravity of pure, 100 per cent sulfuric acid is 
slightly less than that of a 96 to 99 per cent acid, the difference 
being so small that the concentration of the acid cannot be de- 
termined satisfactorily by means of the density. 

When sulfuric acid is mixed with water, considerable heat is 
evolved, and the volume of the diluted acid is considerably less 
than the sum of the volumes of the acid and water which are 
mixed. There is a chemical combination between the acid 
and water, giving a compound which probably contains four or 
six hydroxyl (OH) groups: 


ELECTRON THEORY 181 


se OH 
OH 

O. OH OH } 
No Vou 

Se HOH 0 oq ge HOn > c 

o” Nou a <on 
Ae OH 
OH 

ae 


It is noticeable that while fa epe acid, "Ss¢ So Tks loses 


/ OH O OH 
water easily, sulfuric acid, sk RG dissociates at a much 


OH 
higher temperature and also has a strong tendency to take up 
more water. Along with this strong attraction of the sulfur 
atom for hydroxyl groups, which seems to be so closely con- 
nected with the addition of another oxygen atom, is the fact 
that sulfuric acid is a much stronger acid than sulfurous acid. 
Thus in a “‘ tenth normal’! solution of sulfuric acid about 60 
per cent of the hydrogen is ionized, while in a tenth normal solu- 
tion of sulphurous acid, H2SO3, only about 20 per cent is ionized. 

The Electron Theory. The facts which have just been given 
may be explained, in part, by the electron theory, which has been 
developed rapidly during the last few years. The electron? 
may be defined as an atom of negative electricity. When by 
itself and in rapid motion its mass is approximately one eighteen- 
hundredth of the mass of a hydrogen atom. 

It is supposed that atoms are composed of positive nuclei 
about which electrons are arranged in layers. The nuclei con- 
sist of protons (p. 138) combined with a smaller number of elec- 
trons in such a way that the nucleus has a very small volume 
and a positive charge, balanced by the planetary electrons 
revolving about it. The number of electrons in the outer layer 


1 A solution containing one tenth of a gram atom of replaceable 
eee or one twentieth of a gram molecule of sulfuric acid in one 
iter. 

2 Professor J. J. Thomson uses the name “‘corpuscle” instead of 
- electron, The evidence of the existence of electrons is very positive. 


182 A TEXTBOOK OF CHEMISTRY 


of a neutral atom may vary from one in sodium and the alkali 
metals to 8 in helium, neon and the other “ noble gases.”” When 
there is only one electron in the outer shell, it is supposed that 
this is easily transferred to other atoms and especially to atoms 
of elements like chlorine, which has 7 electrons in its outer shell, 
when it is electrically neutral. The sodium atom which has 
lost an electron in this manner becomes a positive ion, Nat, 
while the chlorine atom, which has gained an electron, becomes 
a negative ion, Cl, in solution. In hydrogen sulfide, when 
completely ionized, the two hydrogen ions, 2 H*, will have lost 
two electrons and the sulfur ion, S, will have gained two. 

In unionized molecules it seems probable that a pair of elec- 
trons furnishes a bond between two atoms. This may be repre- 
sented, graphically, by using dots to represent electrons, thus: 


H: Cl: H:S:H 


In compounds which do not ionize, or which ionize to a slight 
extent, we may suppose that electrons form the bond between 
atoms held in combination. In the case of double unions we 
may suppose that four electrons furnish the bond, thus: 


AO hse ee) Oe 

The condition of the electrons in the bonds between atoms 
is still open to speculation. 

Sulfuric Acid as a Dehydrating Agent. On account of its 
affinity for water, sulfuric acid is an excellent drying agent for 
all gases which do not react with it chemically. It is much more 
efficient than calcium chloride. It will also take the elements 
of water from many such substances as wood or sugar, which 
contain oxygen and hydrogen. 

So much heat is liberated when sulfuric acid is mixed with 
water that the action may become explosive unless care is used. 
The concentrated acid should always be poured into water with 
which it rs to be diluted instead of pouring water upon the acid. 


Why? 


SULFATES. NORMAL SOLUTIONS 183 


Sulfates. Dibasic Acids. [Either one or both of the hydrogen 
atoms of sulfuric acid may be replaced by a metal, giving acid 
and normal salts, as acid sodium sulfate, NaHSOu,, and normal 
sodium sulfate, NagSO.. Acids having this property are called 
dibasic. An acid like phosphoric acid, H3PO., which forms 
three salts with sodium, NaH,POu, NazHPO, and Na3PQu, is 
called tribasic. The basicity depends, however, not on the 
number of hydrogen atoms in one molecule of an acid, but on 
the number of replaceable hydrogen atoms. Thus acetic acid, 
C2H4O2, is monobasic because only one of its hydrogen atoms 
can be replaced ; and phosphorous acid, H3PQs, is dibasic because 
only two of the hydrogen atoms can be replaced. 

As with other strong acids, the normal sulfates of the metals 
of the sodium and calcium families are neutral in reaction, 
while the acid sulfates of all metals are strongly acid. The sul- 
fates of the metals of the calcium family, calcium, strontium, 
barium and radium, are difficultly soluble in water, the solubility 
decreasing with increasing atomic weight. Barium sulfate re- 
quires about 400,000 parts of water for its solution, while radium 
sulfate is still more insoluble. Lead sulfate, also, is almost 
insoluble, but all other sulfates which are not decomposed by 
water are soluble. In general, the salts of strong acids are solu- 
ble in water, and this fact is probably connected with the high 
degree of ionization of both acids and salts. No explanation 
has been offered for the exceptions to this general rule. The 
rule is useful because we have to learn only a short list of insolu- 
ble salts for these strong acids and can then assume that all other 
salts are soluble. 

Normal, Standard and Formular! Solutions. We have fre- 
quently found it convenient to use the gram molecule of sub- 
stances as a unit in dealing with them. This unit is often called, 
for the sake of brevity, one mol. In working with acids and 

1The designation ‘‘molar’’ (or molal) is often used, but ‘‘for- 
mular,’’ if followed by the formula of the substance, is more definite. 
Thus a formular solution of ferric chloride, FeCl;, is definite, while 


2 Pe solution of ferric chloride might refer to either FeCl; or 
Co lg. 


184 A TEXTBOOK OF CHEMISTRY 


bases it is often convenient to use a gram equivalent instead of a 
gram molecule, as the unit. The gram equivalent of an acid 
or base is that quantity which is equivalent to or will neutralize 
one gram molecule of a monobasic acid or of a monacid base. 
A solution containing one gram equivalent of an acid or base in 
one liter (or one milligram equivalent in one cubic centimeter) 
is said to be normal. Thus a normal solution of hydrochloric 
acid would contain 36.47 grams of the acid, HCl, in one liter; 
but a normal solution of sulfuric acid would contain, not a gram 
molecule (98.08 grams), but a gram equivalent (49.04 grams) 
of sulfuric acid, H2SOx, in one liter. A normal solution of sodium 
hydroxide, NaOH, would contain 40.01 grams in one liter; but 
a normal solution of calcium hydroxide, Ca(OH)s, if it could be 
prepared, would contain only 37.04 grams per liter. ‘The advan- 
tage of the system is that one cubic centimeter of any normal 
solution will exactly neutralize or be exactly equivalent to one 
milligram equivalent of any acid or base. The name “‘ normal ” 
is also frequently applied to solutions of salts and of other sub- 
stances, but such a use is liable to lead to confusion; and it is 
better to call such solutions ‘‘ standard” or “ formular,” a 
standard solution being simply one whose concentration 1s 
known and a formular solution one which contains one formular 
weight in one liter. The formula on which a formular solution 
is based should always be given. 

If the term “normal” is applied to solutions of other sub- 
stances than acids and bases, one liter of the solution should 
always contain an amount of the substance which is equivalent 
to one gram atom of hydrogen in the reaction for which it is 
used. A normal solution of potassium chloride, KCl, or of 
silver nitrate, AgNOs, will contain one gram molecule of these | 
compounds in a liter, but a solution of calcium chloride, CaCl, 
will contain only one half of a gram molecule. <A solution of 
potassium permanganate, KMnO,, will contain only one fifth of 
a gram molecule, if to be used in an acid solution, and one third 
of a gram molecule, if to be used in an alkaline solution, because 
one gram molecule of the compound will oxidize five gram atoms 


ACIDIMETRY 185 


of hydrogen in the first case and only three in the second. A 
normal solution of ferrous sulfate, FeSO,, will contain one gram 
molecule of the compound in a liter, because it requires only 
one half of a gram atom of oxygen to oxidize it. 

Acidimetry and Alkalimetry. If a very small quantity of an 
indicator (p. 122) is added to a solution of hydrochloric acid or 
any other strong acid, on adding a 
solution of sodium hydroxide or some 
other strong base, the change in 
color of the indicator will show very 
sharply when the acid has been ex- 
actly neutralized by the base. 

If we have a normal solution of 
hydrochloric acid (containing 36.47 
milligrams of the acid, HCl, in 1 ce., 
as defined above), it is easy by 
measuring this from a burette (Fig. 
55) exactly to neutralize a solution 
containing a strong base. The num- 
ber of cubic centimenters of the acid 
used will give at once the numbers 
of milligram equivalents of the base 
which were present in the solution 
neutralized. Thus one cubic centi- 
meter of normal hydrochloric acid will exactly neutralize 40.01 
milligrams of sodium hydroxide, NaOH, 56.11 milligrams of 
potassium hydroxide, KOH, or 37.04 milligrams of calcium 
hydroxide, Ca(OH):. In the same manner, by means of a 
normal solution of potassium hydroxide, KOH (containing 
56.11 milligrams in 1 cc.), the number of milligram equivalents 
of any strong acid contained in a given solution can be readily 
determined. The process of making such determinations is 
called acidimetry or alkalimetry. The choice of an indicator 
and the application of the process to some cases involving 
weak acids and bases will be discussed in a later chapter 

(p. 387). 


Fig. 55 


186 A TEXTBOOK OF CHEMISTRY 


Pyrosulfates. When acid sodium sulfate is heated, it loses 
water and is converted into a salt which is called, for this reason, 
sodium pyrosulfate : 


Zz NaHSO, = NaeS.O, + H.O 
Sodium 
Pyrosulfate 

A solution of sulfur trioxide in sulfuric acid doubtless always 
contains pyrosulfuric acid, H2S.O;, and pure pyrosulfurie acid 
is a definite compound which melts at 35°, but it is very unstable, 
dissociating easily into sulfur trioxide and sulfuric acid. In 
solution the pyrosulfates take up water and pass back into acid 
sulfates. 

Hyposulfites. When zinc is dissolved in a solution of sulfurous 
acid, H2SOs, the acid is reduced and zinc hyposulfite is formed : 
Zn + 2 HeSO; = ZnS.O,4 + 2 H2O 
Zine 
Hyposulfite 

Zine hyposulfite is a salt of an unstable acid, hyposulfurous 
acid, H2S.O,, which is not known in the free state. The salts 
are very quickly oxidized to sulfites in the air and are powerful 
reducing agents. Sodium hyposulfite, NagS.Ou, is manufactured 
for use in the reduction of indigo to indigo white (p. 341). 

Hyposulfurous acid and the hyposulfites must not be con- 
fused with thiosulfuric acid and thiosulfates which were formerly 
called by the same name (see the next paragraph). Some 
authors prefer to call the acid hydrosulfurous acid to avoid 
possible confusion. Neither acid corresponds to the formula 
(H2SO2), which we should logically expect for a hyposulfurous 
acid. 

Thiosulfates. A solution of sodium sulfite, NasSO3, will 
dissolve sulfur, and there may then be crystallized from the solu- 
tion a salt called sodium thiosulfate, NasS.O3.5 H.O. The 
change is similar to the oxidation of sodium sulfite to sodium 


sulfate: NiceOn Oi Nass 
Na SO; + S = NaeS.O3 


PERSULFURIC ACID 187 


one atom of sulfur taking the place of an atom of oxygen, and the 
name thiosulfate (from Greek 6etov, sulfur) is given to the 
salt for this reason. The salt has been long known and was 
originally called sodium hyposulfite, a name which still clings to 
it among druggists and photographers. It is extensively used 
in photography as a solvent for silver chloride or bromide in 
“fixing ”’ pictures. 

If a solution of a thiosulfate is acidified, the thiosulfuric acid 
at first liberated decomposes with the liberation of sulfur: 


H,.S.03 == SO. of H.O a. S 
Iodine converts sodium thiosulfate into sodium tetrathionate : 


Z NayS2O3 + I, = Na2S40¢ a De Nal 


Sodium 
Tetrathionate 


This reaction is much used in connection with standard iodine 
solutions, in volumetric analysis. 

Persulfuric Acid. When a solution of acid potassium sulfate, 
KHSO,, is electrolyzed with a high current density, that is, with 
a current strong in comparison with the surface, at the anode, 
as the anions, HSO,', are discharged they combine, in part, with 
other anions of the same kind to form persulfuric acid, 
H—O—SO,—O—O—SO,.—OH or H2S:O03. The persulfuric acid 
then reacts with some of the acid potassium sulfate present to 
form potassium persulfate, which is rather difficultly soluble: 


2 KHSO, a H.S.O; 2S K.S0s Z H.SO, 


Persulfuric acid is also formed when hydrogen peroxide, H2Os., 
is added to concentrated sulfuric acid : 


H—O—SO,—|O—-H H'i-O—O—!H H—O'—S0,—0H 


or 2 H.SO, + HO, = H.S2Os 1. 4 H.O 


Persulfuric acid and the persulfates are used as oxidizing 
agents. 


188 A TEXTBOOK OF CHEMISTRY 


* Permonosulfuric Acid. When a solution of persulfuric acid 
is diluted and allowed to stand, it changes to permonosulfuric 
acid : 

H—O—SO,.—-O—O—SO,— OH + HOH 
=H—0O--50,—0— 0H Hs50; 
or H.S820¢ + H.O = H.SO; ++ H.SO, 


The solution was formerly known as Caro’s acid and is some- 
times used as an oxidizing agent for organic compounds. For 
instance, it will oxidize aniline to nitrobenzene. 

* Polythionic Acids. A series of acids having from two to six 
atoms of sulfur in a molecule has been obtained. These are: 

Dithionic acid H.S205 

Trithionic acid H.S30¢ 

Tetrathionic acid H.S4Os. (See thiosulfuric acid, above.) 

Pentathionic acid H2S;O¢ 

Hexathionic acid H2S,0¢ 
These acids need not be considered in detail here. 

Compounds of Sulfur containing Halogens. Quite a num- 
ber of such compounds are known. All of them except sulfur 
hexafluoride, SF¢, are hydrolyzed by water, giving hydrochloric 
acid or a halogen acid and some acid of sulfur. Thus sulfuryl 
chloride, SO2Cle, gives: 


HOH 


* Sulfur Monochloride, S2Cls, is a clear, amber-colored liquid 
formed by passing chlorine over heated sulfur. It boils at 138° 
and is hydrolyzed by water to hydrochloric acid, thiosulfuric 
acid and sulfur. Its specific gravity is 1.7055. 

Because of the strong affinities of sulfur for oxygen and of 
chlorine for metals, the chlorides of a number of metals, which 
it is difficult to prepare from the oxides otherwise, may be ob- 
tained by passing a mixture of sulfur monochloride and chlorine 
over the heated oxides. (E. F. Smith.) It is also used in the 
manufacture of India rubber. 


SELENIUM 189 


* Chlorosulfonic Acid! is easily prepared by passing hydro- 
chloric acid gas through warm, fuming sulfuric acid. 


O Vee 
se) whys 
oe » Nou 


It boils at 152-153° and is easily hydrolyzed by water. Its 
specific gravity is 1.766 at 18°. 

* Sulfuryl Chloride, SO.Cle, may be prepared by the union 
of sulfur dioxide and chlorine, or, more easily, by boiling chloro- 
sulfonic acid with mercuric sulfate, which acts as a catalytic 
agent, causing it to decompose in accordance with the equation : 


2 SOLOHC! = H2SO4 + SO2Cl: 


Sulfuryl chloride is sometimes called, less correctly, the chlo- 
ride of sulfuric acid. Similar acid chlorides may be formed by 
replacing the hydroxyl of other acids with chlorine. Acid chlo- 
rides are hydrolyzed by water, giving the acid from which they 
are derived and hydrochloric acid. 

Sulfuryl chloride boils at 69.1° and has a specific gravity of 


1.6674 at 20-2 


O° . 





Selenium, Se, 79.2, is found in small amounts as selenides of 
metals, associated, usually, with sulfides of these same metals. 
When such sulfides are used for the manufacture of sulfuric 
acid, selenium is sometimes found in the dust flues of the pyrites 
burners and in the slime on the bottom of lead chambers. Both 
selenium and tellurium are found in considerable quantities in 
the slimes from electrolytic copper refining. 

It occurs in several allotropic forms, the red variety obtained 
by crystallization from carbon bisulfide and a gray metallic 
form obtained by melting either of the other forms being the 
best defined. An amorphous form is also known. The metallic 
form melts at 217° and has a specific gravity of 4.8. Selenium 
boils at 680°. 


1 Acids containing the group SO20H are called sulfonic acids. 
* This means the specific gravity at 20° referred to water at 4°. 


190 A TEXTBOOK OF CHEMISTRY 


The metallic form of selenium conducts electricity. Its con- 
ductivity is very greatly affected by changes of temperature or 
by exposure to light, and several important applications of this 
property have been invented, one of the most important being 
in stellar photometry. 

Hydrogen Selenide, H.Se, may be prepared by the action of 
hydrochloric acid on ferrous selenide, FeSe. It is very poison- 
ous, and the odor is more unpleasant than that of hydrogen sul- 
fide. Compare the series, water, hydrogen sulfide, hydrogen 
selenide, in this respect. Berzelius, one of the early workers 
with hydrogen selenide, reports that after breathing a single 
bubble of the gas he so far lost the sense of smell for several 
hours that he could not distinguish the odor of strong ammonia. 

Selenium Dioxide, SeOs, is a white solid prepared by burning 
selenium in a current of oxygen. It gives selenious acid, H2SeOs, 
on solution in hot water. From a solution of selenious acid 
sulfur dioxide precipitates selenium as a red powder: 


H.SeOQ3 -+ 2 SO. ote H,O = Se f+ 2 H.SO,4 


Selenic acid, H2SeQOu, is formed by the action of bromine on 
silver selenite, the silver bromide formed separating as a precipi- 


tate: AgoSeO; + Br. + H.O = H.SeO, + 2 AgBr 


Selenic acid loses oxygen easily and is a strong oxidizing agent. 
Tellurium, Te, 127.5, is found in combination with gold, 
silver, copper and bismuth. It is a white, metallic-looking 
solid, which melts at 452°, boils at 1400° and has a specific 
gravity of 6.44. Its most interesting compounds are hydrogen 
telluride, HeTe, tellurium dioxide, TeOs, tellurium trioxide, 
TeQOs, tellurous acid, H2TeOs3, and telluric acid, H»TeQOu:. 
Atomic Weight of Tellurium. Very many determinations of 
the atomic weight of tellurium have given values about 127.5, 
decidedly higher than the atomic weight of iodine, 126.9. The 
properties of tellurium and especially the formulas of its com- 
pounds indicate that it should precede iodine in the Periodic 
System, and this has led to many attempts to determine whether 


TELLURIUM. GROUP VI 191 


the material used for the atomic weight determinations has been 
pure, or whether, possibly, it may have contained some other 
element from which it is unusually difficult to separate a pure 
tellurium compound. Some of these attempts to discover a 
method of preparing tellurium of greater purity and lower atomic 
weight have seemed, for a time, to be successful; but none of 
these lower results for the atomic weight has been confirmed by 
other workers, and it seems pretty certain that the atomic weight 
of tellurium is greater than that of iodine. See Browning and 
Flint, Z. anorg. Chem. 64, 104, 112, and 68, 251; and Harcourt 
and Baker, J. Chem. Soe. 100, 1311. 

General Properties of the Elements of the Sixth Group. Just 
as chlorine, bromine and iodine are much more closely related 
in their properties than fluorine is related to them, oxygen stands 
somewhat by itself in the sixth group, while the relationships 
between sulfur, selenium and tellurtum are comparatively close. 
The halogens have a valence of one in their compounds with 
hydrogen, as in HF, HCl, ete., and a maximum valence of sevenin 
their compounds with oxygen, as in Cl,0;, HClOs(H—O—ClOs), 
HIQO,, ete. The elements of the sulfur family have a valence of 
two in their compounds with hydrogen, as in H2O, HS, etc., and 
a maximum valence of six toward oxygen, as in SOs, H2SOu,, 


13; 
ys¢ REEO wate 
H-o” o 


There is a similar gradation of physical properties in the two 
groups: fluorine and chlorine are gases, bromine a liquid and 
iodine a solid, with increasing depth of color as the atomic weight 
increases. In the same way, oxygen is colorless (ozone is blue), 
sulfur is a light yellow solid, selenium is dark red, and tellurium 
is opaque and has many of the properties of a metal. Indeed, 
if it were not for its position in the Periodic System and the re- 
semblance between the formulas of its compounds and those of 
selenium and sulfur, tellurium would be classed as a metal, or, 
at least, as a half metal. But it will be seen that in the succes- 

sive groups the metallic properties become more and more 


192 A TEXTBOOK OF CHEMISTRY 


marked with increasing atomic weight. Thus arsenic, antimony 
and bismuth, of the fifth group, are usually classed as metals, 
though all of them are brittle. Tin and lead, of the fourth group, 
are clearly metals and are malleable, though deficient in tenacity. 

The most typical compounds of the sixth group are the follow- 





Spee at tiC) H.S H.Se H.Te 
O3 SO2 SeO2 TeOr 
SO3 ae a TeO; 

= H.SO;3 H.SeO3 H.TeO3 

— H.SO,4 H.SeOQ, H.TeO, 


As in the halogen family, the chemical activity, in general, 
decreases with increasing atomic weight. Hydrogen sulfide 
dissociates at a much lower temperature than water does, and 
sulfur dioxide will take oxygen from selenious acid, reducing 
it to free selenium. 

As manganese forms compounds which resemble some of the 
compounds of chlorine, there are four metals of the sixth group, 
chromium, molybdenum, tungsten (symbol W, from wolfram) 
and uranium, which form oxides and salts of acids similar to the 
oxides and acids of sulfur. The oxides are: CrO3, MoO3, WOs, 
UQO3; and the corresponding sodium salts of the acids are: 
NazCrOu, NazMoO.u, NazWO,. Uranium forms a compound, 
UO,(OH):, similar in composition to sulfuric acid, but it is a 
base rather than an acid, another illustration of the fact that an 
increase in the atomic weight, within a given group, increases 
the metallic properties of the element. 

Crystals. When substances solidify from the molten condi- 
tion or when they separate on the evaporation of a solution, 
molecules of the same kind frequently arrange themselves in 
definite, geometrical relations to each other, forming solids 
bounded by plane faces, which are called crystals. This prop- 
erty has already been mentioned as an important means for 
preparing pure substances. It is also a very important and char- 
acteristic property of individual substances, and the shapes of 
the crystals of different compounds offer such an infinite variety 


CRYSTALS 193 


that they may frequently be used as a very positive means of 
identification. 

In spite of the large number of crystalline forms all crystals 
may be classified in six systems. These systems are most easily 
defined by referring each to axes, which are used in much the 
same manner as the coordinates of analytical geometry, to de- 
fine the structure of the crystal and the relation between the 
planes bounding its surface. 

These systems are: 

1. The Isometric or Regular System. Three equal axes at 
right angles. Some of the simplest forms of this system are the 


f <I re 
Fig. 56 Sigs o7 Fig. 58 


cube (Fig. 56), octahedron (Fig. 57), rhombic dodecahedron (Fig. 
58) and the tetrahedron (Fig. 59). The last has only half of the 


Fig. 59 Fig. 60 








faces of the octahedron and is called a hemihedral form. Com- 
binations of two or more forms are alsocommon. Figure 60 is 
a combination form called a tetrahexahedron. 


194 A TEXTBOOK OF CHEMISTRY 


2. The Tetragonal System. Three axes at right angles, two 
of them, only, being equal. The tetragonal pyramid (Fig. 61) 
and the square prism (Fig. 62) are common forms. Figure 63 
shows a combination of the two. 





Fig. 61 


3. The Rhombic System. Three axes at right angles but of 
unequal length. Rectangular and rhombic prisms (Figs. 64 and 
65) and pyramids are illustrations. 





Fig. 64 Fig. 65 


4. The Hexagonal System. Three axes in the same plane at 
an angle of 60° with each other and a fourth at right angles to 


CRYSTALS 195 


the plane of the other three. Common forms are the hexagonal 
pyramid (Fig. 66) and prism (Fig. 67) and the rhombic hexahe- 
dron (Fig. 68), which has only half of the faces of the pyramid. 


\ 
ek, 
x] 


Fig. 66 Fig. 67 





5. The Monoclinic System. ‘T'wo axes at right angles and a 
third at right angles to one and inclined to the other, the three 
axes being unequal (Fig. 69). 





Fig. 68 Fig. 69 


WNIVERSITY OF ILLINOIS 
LIBRARY-CHEMIST XY 


196 A TEXTBOOK OF CHEMISTRY 


6. The Triclinic System. Three unequal axes, all inclined 
(Fig. 70). : 

Crystals rarely exhibit the complete geometrical forms which 
are the ideal to which they are referred. They always have an 
internal structure characteris- 
tic of these forms, however, 
and this can often be detected 
by their optical properties. 
Thus for the regular system 
light travels with the same 
velocity in all directions 
through the crystal. For 
other, less symmetrical forms, 





Fie. 70 Serer the velocity is different in 

different directions and this 

causes double refraction, polarization of light and other phe- 

nomena frequently used for the identification of the crystal form. 

The angles between the faces of crystals are also accurately 

fixed by the system to which they belong and the properties 

of the individual substance, and the measurement of these angles 
is used for purposes of identification. | 

Such substances as sulfur, which crystallize in two different . 
forms, are called dimorphous. 

Different substances which crystallize in the same form are 
called isomorphous. A crystal of a substance should grow if 
placed in a supersaturated solution of a substance with which 
it is isomorphous. 


EXERCISES 


1. One liter of water at 0° absorbs 4.37 volumes of hydrogen sul- 
fide. What part of a gram molecule of the gas does the solution con- 
tain? What would be the theoretical depression of the freezing point 
if the compound were completely ionized to HS~ and H*? If it were 
not at all ionized ? 


2. How many liters of air containing 21 per cent of oxygen will be 
required to burn one liter of hydrogen sulfide to water and sulfur ? How 
many liters would be required to burn it to sulfur dioxide and water ? 


SULFUR, SELENIUM AND TELLURIUM 197 


3. How many cubic feet of air at 0° (1 cu. ft. = 28.315 liters) will be 
required to burn enough pyrites to make one ton of chamber acid of 
70 per cent? (1 ton = 907.18 kilograms.) How many cubic feet of 
air must be introduced into the chamber to convert the sulfur dioxide 
to sulfuric acid ? 


4. Solve the same problem, substituting the metric ton and cubic 
meters for ton and cubic feet. How many kilograms of water must be 
introduced in the chamber ? 


5. What is the percentage increase in the volume of the air which 
must be used if the temperature is 25° instead of 0°? 


6. How many grams of ferric oxide will be obtained by burning one 


kilogram of iron pyrites ? 


CHAPTER XII 


NITROGEN 


SymsBou, N. Atomic Weiaut, 14.01. 


Occurrence and Natural History of Nitrogen. Fluorine and 
oxygen, the first elements of the seventh and sixth groups, are 
very active; and while oxygen is found free in the air, this seems 
to be more on account of its abundance and because nearly all of 
the other elements in nature are already combined with oxygen 
than because of any lack of activity. Nitrogen, the first element 
of the fifth group, in very striking contrast to oxygen and fluorine, 
is found chiefly in the free state in the atmosphere because it 
does not readily combine with any of the other elements avail- 
able in the earth. 

Until 1894 it was supposed that the gas remaining when oxy- 
gen, carbon dioxide and moisture were removed from air was 
pure nitrogen. Rayleigh and Ramsay showed at that time that 
the residue left after the removal of these substances still con- 
tained about 1.2 per cent of gases, chiefly argon (p. 235), which 
are even more inert than nitrogen. Nitrogen forms, however, 
nearly 78 per cent of the volume of dry air; and as the air above 
a square centimeter of the earth’s surface weighs about one 
kilogram or that above a square meter weighs more than ten 
tons, it is evident that the amount of nitrogen in the atmosphere 
is very large, though it is small in comparison with the amounts 
of those elements which make up the bulk of the solid crust 
of the earth. In combination, nitrogen is found in all living 
organisms as an essential constituent, the four most important 
elements in organic matter being carbon, nitrogen, hydrogen and 
oxygen. But while plants can obtain the carbon for their growth 
from the carbon dioxide of the air and the oxygen and hydrogen 
from the moisture of the soil, very few, if any, of the higher 


forms of plant life can use the nitrogen of the air directly. The 
198 


NITROGEN 199 


larger part of the nitrogen which is essential for the growth of 
crops must be supplied in the form of compounds which result 
from the decay of animal or vegetable substances, or from the 
combination of oxygen and nitrogen in the air through electrical 
agencies. There are, however, a few plants, especially clover, 
alfalfa and other leguminous plants, which are able to assimilate 
the nitrogen of the air with the aid of bacteria which grow in 
nodules on their roots. 

The decay of organic matter containing nitrogen is always 
caused by the growth of bacteria. In the absence of air, the 
conditions of decomposition lead to the reduction of the nitrogen 
to ammonia. In the presence of air, as in a well-aérated soil, 
the nitrifying bacteria, which are usually present, will convert 
the nitrogen to nitric acid, HNO;, which generally finds enough 
potassium, calcium or sodium present to form saltpeter, KNOs, 
calcium nitrate, Ca(NO3)2 or sodium nitrate, NaNO;. All 
plants can readily assimilate the nitrogen of the nitrates and so 
it finds its way back into the organic compounds of the plant 
life. As nitrogen in a readily available form is essential to the 
growth of wheat, corn and other crops, sodium nitrate, the 
cheapest of the commercial nitrates, and ammonium sulfate, 
(NH,)2SOu, also a comparatively cheap nitrogen compound, are 
often used as fertilizing materials. 

The course of nitrogen in nature is shown by the following 
diagram :! 

Leguminous plants with 





the help of bacteria 










Atmospheric 
electricity Plants 
Atmospheric | ——————> |... Organic compounds 
; P Nitrates |——> 8 : P 
BOE Clg eee mee of Nitrogen 
Denitrifying Niftrifying Decayjof plant 
bacteria 


bac|teria and anijmal tissues 


A or distillation 
Nitrifying ; 
Nitrites |}<—— | Ammonia 
bacteria Ammonia, 


1 See Abegg, Handbuch der anorg. Chemie, Bd. 3, Abth. 3, S. 215. 





200 A TEXTBOOK OF CHEMISTRY 


Preparation and Properties of Nitrogen. Nitrogen which is 
pure with the exception of about 1.2 per cent of the inert gases 
of the argon family may be prepared by burning phosphorus 
in air or by passing air over heated copper turnings, which 
will take up the oxygen. A very convenient method for the 
preparation of considerable quantities of nitrogen is to pass a 
mixture of air and hydrogen through a tube containing hot 
copper oxide. By keeping the hydrogen slightly in excess of 
the amount necessary to combine with the oxygen, a part of the 
copper oxide will be reduced to metallic copper; and it is then 
easy to regulate the currents of the two gases so that both 
copper and copper oxide will be present in the tube. The hy- 
drogen must, of course, be mixed with the air at the point where 
the two gases come in contact with the copper. (Why ?) 

To prepare nitrogen free from argon some compound of nitro- 
gen must be used, ammonium nitrite, NHsNOs, being most 
suitable. When a solution of the salt is warmed, it decomposes 
to water and nitrogen : 


NH,NO, = No + 2 H.O 


Instead of ammonium nitrite a mixture of sodium nitrite and 
ammonium chloride may be used, the following reversible reac- 
tion occurring first : 


NaNO, + NH.Cl = NHiNO: + NaCl 


Nitrogen is a colorless, odorless gas which condenses to a 
liquid that boils at — 196° and freezes at — 210°. It is very 
inert and does not combine with any element at ordinary tem- 
peratures, except under the influence of micro6rganisms, as 
referred to above. At the high temperatures of the electric 
discharge it combines with oxygen to form nitric oxide, NO, 
and this fact has been used recently as a basis for the commercial 
manufacture of nitrates (see below). As the supply of sodium 
nitrate in Chile, the only large supply now known, will be ex- 
hausted within a comparatively few years, it seems certain that 
this manufacture is destined to be very important. 


AMMONIA 201 


At moderate temperatures and under high pressure nitrogen 
and hydrogen combine to form ammonia, NH3, but the reaction 
is very slow without some catalytic agent. 

The combination is exothermic and hence the equilibrium, 


N: + 3 Hex 2 NH; 


is shifted toward the decomposition of the ammonia at high 
temperatures. As the volume decreases as the gases combine, 
pressure shifts the equilibrium to the right. Principle of van’t 
Hoff-Le Chatelier, p. 111. (See Haber, Z. Elektrochem. 16, 
244.) The study of the conditions for the successful synthesis of 
ammonia advanced so far before 1914 that nitric acid obtained by 
the catalytic oxidation of ammonia became the chief source for 
the explosives used by Germany during the war. Without this 
process, Germany would have been quickly defeated. Important 
modifications of the process have now been developed in France, 
England and America. For the cyanamide process, see p. 462. 
Several metals, especially lithium, magnesium or calcium, 
combine with nitrogen to form nitrides, at high temperatures: 





| yu 
6Li +N, =2Li,N | or NCLi 
Lithium Ni; 

Nitride 

( ) 

we 

3Mg + No = Mg3No_ | or Me | 

okie 
Nine” | N\\,,| 


Nitrogen will not support combustion nor burn. 

Ammonia. When organic matter containing nitrogen decom~- 
poses with exclusion of air, either under the influence of bacteria 
or of heat, the nitrogen is converted partly into ammonia, NHs3. 
In this way ammonia is always found in sewage or in piles of 
manure. It is also formed in the destructive distillation of bi- 
tuminous coal for the manufacture of illuminating gas. The 
aqueous portion of the liquid distillate from the coal furnishes, 
at present, the chief source of the ammonia of commerce. These 


202 A TEXTBOOK OF CHEMISTRY 


ammontacal gas liquors are mixed with slaked lime and distilled, 
the lime retaining sulfur and other impurities. The distillate 
is mixed with hydrochloric or sulfuric acid and evaporated to 
obtain ammonium chloride, NH«4Cl, (NH3 + HCl), or ammo- 
nium sulfate, (NH4)2SO.u, (2 NH3 + H2SQ,). 

From these salts the ammonia may be liberated by any strong 
base, as sodium hydroxide or calcium hydroxide: 


(NH,)2SO4-+ Ca(OH)22% CaSO.+ 2 NH.OH = 2 NH;+ 2 HO 
Ammonia may also be obtained by hydrolyzing a nitride with 
water : 


LijsN + 3 HOH = NH; + 3 LiOH 


For laboratory or lecture purposes ammonia gas is most easily 
obtained by boiling a strong solution known as aqua ammonia and 
passing the gas through a cylinder filled with quicklime to dry it. 

Properties of Ammonia. Ammonia is a colorless gas with a 
very pungent odor. It is very easily soluble in water, and hence 
must be collected by displacement of air (should the mouth of 
the bottle point up or down?) or over mercury. Water at 0° 
absorbs about 1000 times its volume of the gas, but gives off a 
large part of it on warming gently and all of it on boiling. The 
density of the solution is less than that of water, a 28 per cent 

solution having a specific gravity of 0.90. 

Ammonia combines directly with acids to form ammonium 
salts, in which the hydrogen of the acid combines with the am- 
monia to form the ammonium group, NH4g, a radical which in 
its compounds possesses properties very closely resembling the 
properties of potassium or sodium : 

NH; + HCl = NH.Cl 


Ammonium 
Chloride 


2 NH; + H.SO, = (NH4)2SO4 
Ammonium 
Sulfate 


These compounds are most satisfactorily explained by sup- 
posing that nitrogen is trivalent when combined exclusively with 


AMMONIA 203 


hydrogen or with positive groups but may become quinquivalent 
when one of the groups or atoms is negative: 


Cl 


Any acid may take up as many molecules of ammonia as it 
has of replaceable hydrogen atoms. Thus a monobasic acid, 
as nitric acid, HNO;, may combine with one molecule of am- 
monia, forming ammonium nitrate, NH,NOs3, or a tribasic 
acid, as H3;PQO., may combine with three molecules. 

The formation of an ammonium salt may be very prettily 
illustrated by filling two cylinders with ammonia and hydro- 
chloric acid gas respectively. On bringing the mouths of the 
cylinders together the gases will combine to form solid am- 
monium chloride. 

Aqua Ammonia. The solubility of ammonia in water has 
already been mentioned. It dissolves, in part, without chemical 
change, as is shown by the strong odor of the solution due to 
the escape of the gas, but it partly combines with the water, 
which easily separates into H* and OH’, forming ammonium 
hydroxide: 


OH 


This ionizes to form ammonium, NH,t, and hydroxide, OH’, 
ions, but the ionization is small in comparison with that of 
strong bases. A tenth normal solution of sodium hydroxide is 
ionized to the extent of about 84 per cent, while a tenth normal 
solution of ammonium hydroxide (or ammonia) shows only 1.3 
per cent of ionization, if we assume that all of the ammonia in 
the solution has combined with water to form ammonium 


204. A TEXTBOOK OF CHEMISTRY 


hydroxide. A normal solution, on the same basis, shows an 
ionization of only 0.3 to 0.4 per cent. How may the presence 
of hydroxide ions in a solution of ammonium hydroxide be 
demonstrated ? 

It is probable, however, that a large part of the ammonia 
exists as such in the solution.! In other words, the solution 
contains ammonia, NH;, as well as ammonium hydroxide, 
NH,OH, and ammonium, NH,*, and hydroxide, OH, ions. 
The practical effect, that ammonium hydroxide is a weak base 
because its solution contains few hydroxide ions, is the same 
whether we suppose this to be because ammonium hydroxide 
is only slightly ionized or whether it is because the ammonium 
hydroxide is largely dissociated into ammonia and water. 

Ice Machines. Ammonia may be readily condensed to a 
liquid either by pressure (4.19 atmospheres at 0°) or by cold 
(— 33° at 760 mm.). The heat of vaporization of the liquid 
is 330 calories per kilogram (at — 33°). This high value is 
intimately connected with the low molecular weight and also 
with the fact that liquid ammonia is, like water, a highly “ asso- 
ciated ”’ liquid, that is, consists of polymerized molecules such 
as (NH3)>. or (NH3)3.. Much heat is absorbed in the vaporiza- 
tion both because of the large volume of the vapor in proportion 
to its weight and because the polymerized molecules must be 
broken up. The high heat of vaporization is utilized in ice 
machines. The principle of one form of these machines is illus- 
trated in the diagram (Fig. 71). Liquefied anhydrous ammonia 
is allowed to evaporate in the coil A and the escaping gas 1s com- 
pressed by the pump B and condenses to a liquid in the coil 
C. From this coil the liquid ammonia is returned through the 

1 Moore, J., Chem. Soc. 91, 1882 (1907), calculates on the basis 
of the partition of ammonia between water and chloroform at dif- 
ferent temperatures that from 30 to 40 per cent of the ammonia 
exists as ammonium hydroxide, NH,OH, at 20°. This does not 
seem to be consistent, however, with the relative ionization constants 
of trimethyl amine, (CH;)3N, and tetramethyl ammonium hydroxide, 
(CH;)4 NOH, which indicate strongly that the former forms only a 
small amount of trimethyl ammonium hydroxide, (CH3);NHOH, in 


aqueous solutions. Further experimental evidence on this ques- 
tion seems highly desirable. 


DERIVATIVES OF AMMONIA 205 


regulating valve D to the coil A. The coil Cis surrounded by 
cold water to absorb the heat evolved as the ammonia condenses. 
The coil A is surrounded with brine, either a solution of salt 
or of calcium chloride. The cold brine may be circulated by 
means of pumps through coils of pipe in refrigerator rooms, or 
cases of distilled water may be immersed in the brine, to be 
frozen. 

In another form of machine, which was in earlier use, the 
pressure to condense the ammonia was obtained by heating a 





Fig. 71 


concentrated aqueous solution. Afterwards the weakened 
solution was cooled, and as it reabsorbed the gas the evapora- 
tion of the liquefied gas caused the refrigeration. 

* Derivatives of Ammonia. One or more atoms of hydrogen 
in ammonia may be replaced by a metal, giving such compounds 
as sodium amide, NaNHz, or by radicals, especially by organic 
radicals, giving such compounds as methyl amine, CH;NH2, 
phenyl amine or aniline, CsH;NHe, acetamide, C2H;ONHe, and 


CO 
phthalimide, CoH NE. If the group replacing hydrogen 
CO 


is a hydrocarbon radical, as methyl, CHs, or phenyl, CeHs, the 
compound is called an amine. The amines combine with acids, 
as ammonia does, to form such salts as methyl ammonium 
chloride, CH3NH3Cl (or CH3NH2.HCl) and phenyl ammonium 
chloride, CsH;NH3Cl. For this reason the amines are often 
called organic bases, but just as ammonia is a base only when it 


206 A TEXTBOOK OF CHEMISTRY 


has combined with water to form ammonium hydroxide, NH,OH, 
the amines are true bases only when combined with water. As 
ammonium hydroxide dissociates to ammonia and water even 
in solution, such hydroxides as methyl ammonium hydroxide, 
CH;NH;0H, dissociate readily into water and the original 
amine and can exist as pure compounds only at very low tempera- 
tures, if at all. All four of the hydrogen atoms in the ammo- 
nium group, NHgz, of ammonium hydroxide may be replaced by 
hydrocarbon radicals, however, and some compounds formed in 
this manner no longer dissociate when their solutions are evapo- 
rated. Thus a white crystalline mass, doubtless consisting of 
tetramethyl ammonium hydroxide, (CH;)s4NOH, separates on 
evaporating a solution of this compound. ‘The preparation of 
this and several other similar compounds has been one of the 
reasons for believing that solutions of ammonia in water contain 
ammonium hydroxide, NH,OH. 

If the group replacing a hydrogen atom in ammonia is an 
acid radical, the compound is called an amide. Thus the com- 
pound containing the acetyl group, C2H;0, is called acetamide, 
CsH3;O0NHe. In aqueous solutions the amides are usually am- 
photeric; that is, they have both very weak acid and very weak 
basic properties. When two hydrogen atoms of ammonia have 
been replaced by a bivalent acid radical, the compound is 


Hons 
called an imide, as phthalimide, C.H, NH. In the 
SHE 


imides the hydrogen can be replaced by metals forming 


The Electron Theory. A comparison of the compounds oi 
nitrogen on the basis of the electron theory (p. 181) is suggestive. 


JY 
In nitric acid, H-O-N€ , the nitrogen atom gives five 
O 


ELECTRON THEORY 207 
electrons to the oxygen atoms, forming a compound which 
readily ionizes to hydrogen, H*, and _ nitrate, Om DER 

On 


ions because of the strongly positive nitrogen atom, which holds 
the oxygen of the hydroxyl firmly but repels its hydrogen. In 


H 
ammonia, H-NC , the nitrogen atom receives three elec- 
H 


trons from the hydrogen atoms, becoming negative. It can 
receive a fourth electron from another hydrogen atom only in 
case it also gives up one electron to the oxygen of a hydroxyl 
group or to chlorine or some other negative atom or group. In 

H* H 
the ammonium hydroxide, HN 4 , which results, 
in the first case, the negative nitrogen atom no longer holds the 
negative oxygen of the hydroxyl group strongly, and so the com- 
pound may ionize toammonium, NH,*, and hydroxide, OH, ions. 
If one of the hydrogen atoms of ammonia is replaced by a nega- 
tive radical, as acetyl, C2H;0, giving the compound acetamide, 
CH , the nitrogen atom no longer takes up hydrogen 


HN 

Ww 
and hydroxyl readily to form a base, or the elements of an 
acid to form a salt. This seems to be because the presence of 


the negative acetyl group, C2H;0, so far reduces the positive 
character of the group PONS ae that it cannot form 
H-SN 


Wy’ 
the positive ion of a salt. 

* Solutions in Liquid Ammonia. Anhydrous ammonia may 
be condensed to a liquid which boils at — 33.5°. This liquid 
ammonia dissolves many substances, and the conductivity of 
the solutions indicates that some of these ionize in the ammonia 
as acids, bases and salts ionize in solutions in water. While a 
very large portion of our study of chemistry deals with reactions 


208 A TEXTBOOK OF CHEMISTRY 


in aqueous solutions, there are closely parallel phenomena in 
ammoniacal solutions. In such solutions we may consider 
ammonia as consisting of H and NHs just as we think of water as 
consisting of H and OH. As in aqueous solutions derivatives 
of water ionize to form hydrogen, H*, or hydroxide, OH’, ions 
according to the nature of the radical, so in ammonia, com- 
pounds which are derivatives of ammonia may ionize to form hy- 
drogen, H*, or amide, NH2°, ions. Thus acetamide, 
C;:H3;0NHsg, ionizes to C2H;0NH™7 and Ht in solution 
in ammonia and is to be considered as an acid in sucha 
solution. Sodium amide, NaNHg, on the other hand, 
ionizes to sodium, Nat, and amide, NH’, ions and is to 
be considered as a base. Curiously enough the latter will 
cause phenolphthalein to turn red in the ammonia solu- 
tion just as hydroxide ions cause it to turn red in aqueous 
solutions. Neutralization in such a solution must con- 
sist in the union of hydrogen, H*, and amide, NHe", ions 
to form ammonia. See Franklin and his coworkers, 
Am..ChemsJ..20,.820 and $26 3215:38" (23. 27 fee couooe 
47, 285; J. Am. Chem. Soc., 26, 499; 27, 192, 820. 
The Volumetric Composition of Ammonia. The ratio 
between nitrogen and hydrogen in ammonia may be 
demonstrated by filling a tube, Fig. 72, with chlorine 
gas, allowing a small amount of concentrated aqua am- 
monia to enter it and following this with some dilute 
sulfuric acid. On allowing water to enter the tube till 
the gas remaining in it is at atmospheric pressure, the 
Fig. 72 tube will be found to be one third full of nitrogen. 
Under the conditions of the experiment the tube full of 
chlorine, which we know is capable of combining with its own 
volume of hydrogen, takes this amount of hydrogen from the 
ammonia and liberates the equivalent amount of nitrogen. In 
other words, one volume of nitrogen is combined with three 
volumes of hydrogen. 
As the gram molecular volume of ammonia weighs 17 grams, 
the complete reaction is : 


COMPOSITION OF AMMONIA 209 





The primary reaction between ammonia and chlorine is : 
3 NH; + 6 Cle = NCI; + Ne + 9 HCl 
This is followed by the reaction : 
NCl; + NH; = Ne + 3 HCl 


which is favored in an acid solution.! The final result is the 
same as though the simple reaction 


took place. As in this reaction three molecules of chlorine, Clo, 
liberate one molecule of nitrogen, No, it is evident that in accord- 
ance with Avogadro’s law a tube full of chlorine should liberate 
from ammonia one third of a tube full of nitrogen. . 

The composition of ammonia by volume may also be shown 
by another, quite different, experiment. If a small amount of 
ammonia gas is introduced into the apparatus shown in Fig. 73, 
and electric sparks are passed between the platinum wires, which 
pass through the walls of the tube, the ammonia will be decom- 
posed into hydrogen and nitrogen. The reaction is reversible, 
but with the equilibrium very far toward the side of decomposi- 
tion at the temperature of the electric discharge: 


2NH;*5Ne+ 3 He 


As two molecules of ammonia give one molecule of nitrogen 
and three molecules of hydrogen, the volume of the gas would 
be doubled when the decomposition was complete. 


1 J. Am. Chem. Soc. 23, 460. 


210 A TEXTBOOK OF CHEMISTRY 


On the other hand, if a mixture of one volume of nitrogen with 
_ three volumes of hydrogen is placed in the same apparatus and 
a little dilute sulfuric acid is introduced, on passing electric 
sparks through the mixture as before, the nitrogen and hydrogen 
will slowly combine, and, as the ammonia 
formed will be absorbed by the sulfuric acid, 
the combination may be carried to comple- 
tion in spite of the unfavorable character of 
the equilibrium. 

Nitric Acid. The two compounds, ammonia 
and nitric acid, are to be considered as the 
fundamental ones for nitrogen. All other com- 
pounds of the element tend to return to one or 
the other of these, or their salts, or else to de- 
compose with the liberation of free nitrogen. 
And all compounds of nitrogen prepared in the 
laboratory, except those derived from organic 
materials, are prepared directly or indirectly 
from ammonia or nitric acid. It seems best, 
therefore, to speak of nitric acid next, though 
such an order of treatment is practical rather 
than logical. 

The formation of nitrates in the soil by the 
action of nitrifying organisms has been referred 
to. The present commercial source for nitrates 
is almost exclusively the sodium nitrate, NaNOs, 
or Chile saltpeter, found in enormous beds in Chile, in South 
America. From this nitric acid is prepared by a process similar 
to that for the preparation of hydrochloric acid. Nitric acid isa 
stronger acid than sulfuric, but if sodium nitrate is mixed with sul- 
furic acid and the mixture heated, the equilibrium of the reaction, 


NaNO; -f- H.SO, eee NaHSO, -}- HNO; 





Piga(3s 


is displaced to the right as the nitric acid distills away from the 
mixture, the boiling point of nitric acid being much lower than 
that of the sulfuric acid. Nitric acid is not, however, very stable, 


NITRIC ACID PA 


and part of it decomposes, forming oxygen, water and nitrogen 
peroxide, NOz, when the distillation is under atmospheric pres- 
sure. To avoid this the operation is sometimes carried out at a 
lower temperature by reducing the pressure. 

Pure nitric acid is a colorless liquid, which boils at 86° and 
has a specific gravity of 1.52. The addition of water causes a 
rise in the boiling point, the acid of maximum boiling point, 
120°, containing 66 to 70 per cent of pure acid. 

Nitric acid is a strong acid, the tenth normal solution being 
ionized to the extent of 92 per cent, while a tenth normal solution 
of hydrochloric acid is 91 per cent ionized. 

* Hydrates of Nitric Acid. The addition of increasing amounts 
of nitric acid to water lowers the freezing point till an acid con- 
taining 32.8 per cent of nitric acid freezes at — 43°. Further 
addition of nitric acid causes the freezing point to rise and fall 









































TEMPERATURE OF FREEZING. 




















0 10 20 30 40 50 60 70 80 90 100% 
PER CENT. OF NITRIC ACID. HNO; 
Fig. 74 


alternately as shown in Fig. 74, pure nitric acid freezing at 
— 41.2°. The two maximum freezing points shown in the figure 
correspond to acids containing 53.84 and 77.77 per cent of 


212 A TEXTBOOK OF CHEMISTRY 


nitric acid respectively. In accordance with the principle that 
a pure liquid freezes at a higher temperature than when it con- 
tains some dissolved substance, these melting points indicate 
definite compounds. An acid containing 53.84 per cent of nitric 
acid corresponds to the formula HNO3.3 H.O and one contain- 
ing 77.77 per cent, to the formula HNO;.H2O. The freezing 
point curve demonstrates very clearly the existence of these two 
hydrates.! 

Chemical Properties of Nitric Acid. As an acid, nitric acid 
has the usual properties, forming salts as a monobasic acid with 
practically all metals. ‘These may be prepared by the action of 
the acid on the metal or on a hydroxide, oxide or carbonate of 
the metal : 

NaOH + HNO; = NaNO; + HOH 
Ca(OH). + 2 HNO; = Ca(NO3)o + 2 H.O 
ZnO ot 2 HNO; = Zn(NOs3)o aa H.O 


The most important special properties of nitric acid depend 
on the ease with which it gives up oxygen to a great variety of 
substances. It is a powerful oxidizing agent, the concentrated 
or the anhydrous acid, HNOs, showing this property in a more 
marked degree than the dilute acid. Indeed the pure acid de- 
composes spontaneously into water, oxygen and nitrogen diox- 
ide, NOs, on distillation or on exposure to light: 


2 HNO; =.2.NO.--.0. 20 


If pure nitric acid is boiled in a test tube having a plug of wool 
or feathers in its mouth, the latter will catch fire and burn. 
Ignited charcoal will also continue to burn beneath the surface 
of the liquid. 

When nitric acid acts upon a metal, hydrogen is rarely, if ever, 
liberated. Instead of this the nitric acid is reduced, either by 
the metal directly, forming an oxide of the metal, or by the hy- 
drogen displaced by the metal. Such hydrogen is often called 


1 Kiister u. Kremann, Z. anorg. Chem. 4/7, 1 (1904). 


AQUA REGIA 213 


. 


“nascent”?! hydrogen. The two ways of explaining the action 
may be illustrated as follows : 


3 Cu + 2 HNO; = (8 CuO) + HO + 2 NO 
Or 3 Cu + 6 HNO; = 3 Cu(NOs;); + (6 H) 
(6 H) + 2 HNO; = 4 H,O + 2 NO 











The final reaction is the same whichever explanation of the 
mechanism of the reaction is adopted. 

Nitric acid may be reduced less or more by other metals, or by 
copper when the concentration of the nitric acid is different, 
giving the whole series of oxides, NOx, N2O3, NO and N.O, and 
even ammonia, NH;3. In general a concentrated acid gives the 
higher oxides, while a more dilute acid or a metal which has a 
high heat of oxidation gives the lower oxides or ammonia. If 
nitricacid is added to a test tube containing zinc and sulfuric acid, 
the evolution of hydrogen may nearly cease and ammonium 
sulfate will be formed. The addition of an excess of sodium 
hydroxide to the solution will then liberate ammonia, which may 
be recognized by its odor and effect on litmus paper. What is 
the series of reactions involved in the experiment ? 

Aqua Regia. Neither nitric acid or hydrochloric acid alone 
will dissolve the so-called noble metals, gold or platinum, but a 
mixture of the two will dissolve them readily. Such a mixture 
is called aqua regia? because of this property. The nitric acid 
acts upon hydrochloric acid in the same manner as other oxidiz- 
ing agents, liberating chlorine, and this attacks the gold or plati- 
num, forming soluble chlorides. 

Because it contains chlorine and oxides of nitrogen, aqua regia 


1 Nascent means in the state of being born. The theory is that 
hydrogen atoms when first liberated are more active than the same 
atoms when they have combined with others to form hydrogen 
molecules, H2. The activity is dependent largely on the metal used, 
however, metals which dissolve with a large evolution of heat giving 
the most active hydrogen. 2 Royal water. 


— 


214 A TEXTBOOK OF CHEMISTRY 


is a powerful oxidizing agent and is frequently used for that pur- 
pose, especially to oxidize the sulfur of sulfides to sulfuric acid, 
as in the determination of sulfur in iron. 

Nitrosyl chloride, NOCI, a volatile compound with a very 
disagreeable odor, is also formed in the mixture of nitric and 
hydrochloric acids. This is hydrolyzed by water in the same 
manner as other nonmetallic chlorides: 


NOCI + HOH = NOOH (or HNO) + HCl 


Nitrosyl chloride is also a strong oxidizing agent. 

Oxides of Nitrogen. There are six oxides of nitrogen, but as 
two of these, nitrogen peroxide, NOs, and nitrogen. tetroxide, 
N.O,, have the same percentage composition and change each 
into the other on merely changing the temperature, they are 
frequently spoken of as a single substance and the name nitrogen 
peroxide is applied to both. The oxides are: 


Nitrous oxide NO 
Nitric oxide NO 
Nitrous anhydride N.O3 
Nitrogen dioxide NO, | 
Nitrogen tetroxide N2O,} 
Nitric anhydride  N.O; 


Nitrous oxide, N2O, is most easily prepared by heating am- 
monium nitrate; nitric oxide, NO, by the reduction of nitric 
acid with metallic copper; nitrous anhydride, NOs, by reducing 
nitric acid with arsenious oxide, As2O3; nitrogen dioxide, NOs, 
by heating lead nitrate, Pb(NO3)2; and nitric anhydride, N2Os,, 
by dehydrating nitric acid with phosphoric anhydride. Nitric 
oxide is also formed by the direct union of the elements in an 





electric arc, and it unites with oxygen spontaneously to form 
nitrous anhydride, N.O3, and nitrogen dioxide, NOs. 

Nitrous Oxide, N.,O. When ammonium nitrate, NH4NQOs, 
is heated, the hydrogen of one part of the molecule combines 
with oxygen from another part to form water, while the two 
nitrogen atoms remain combined with the other oxygen atom : 


NH4NO; = NO + 2 HO 


NITRIC OXIDE Zio 


The reaction is exothermic, that is, it takes place with evolu- 
tion of heat and is hable to become explosive if the temperature 
is raised too high or if too large a quantity of the salt is heated 
at once. 

Nitrous oxide is a colorless gas with a sweetish odor and taste. 
Water at 20° absorbs about two thirds of its volume of the gas. 
It supports combustion. A glowing splinter will inflame in 
the gas somewhat as it does in oxygen, and phosphorus burns in 
it with an intense light. It will not support life. 

When inhaled, nitrous oxide sometimes causes hysterical laugh- 
ing, and it is called for that reason laughing gas. In larger 
amounts it produces insensibility and is used for this purpose 
in minor surgical operations, especially for the extraction of 
teeth. 

The structure of nitrous oxide is probably that represented by 


N 
the formula, | 0. This accounts best for the ease with which 
N 


it gives up oxygen and reverts to free nitrogen. It is the only 
oxide of nitrogen in which two atoms of nitrogen are supposed 
to be directly united. 

Nitric Oxide, NO, is formed when copper dissolves in dilute 
nitric acid of specific gravity 1.2. Ifa stronger acid is used, some 
nitrogen dioxide, NOs, will be formed; while if the acid is much 
more dilute, nitrous oxide, NO, and nitrogen, Ne, will be formed 
along with the nitric oxide. The mechanism of the reaction has: 
already been discussed. 

Nitric oxide, NO, is also formed by the direct union of nitrogen 
and oxygen in an electric arc. For instance, if electric sparks 
from an induction coil are passed for some time between terminals 
in a large globe, the air in the globe will gradually become reddish 
brown in color from the formation of nitric oxide, which combines 
with more oxygen to form nitrogen dioxide, NO». There is evi- 
dence that the combination is caused by the high temperature 
of the are and not by the electricity as such. 

The combination of nitrogen and oxygen is an endothermic 


216 A TEXTBOOK OF CHEMISTRY 


reaction; that is, heat is absorbed as it proceeds. It is also a 
reversible reaction : 

O. + Nez 2 NO 
with the equilibrium very far to the left, so far, indeed, that the 
amount of nitric oxide formed from the elements is very small, 
even at high temperatures. The per cent of nitric oxide formed 
in air when the reaction comes to equilibrium is as follows : 4 














ABSOLUTE PER CENT oF NO 


TEMPERATURE CALCULATED OBSERVED 
Isis 0.35 0.37 
2195° 0.98 0.97 
2675° Da | 9.93 
3200° 4.43 About 5. 








With the equilibrium so far on the side toward its decomposi- 
tion, it seems at first thought that nitric oxide ought not to exist 
at all at ordinary temperatures, and it could not except for the 
fact that the speed of the formation or decomposition is very 
slow. Thus it has been shown that for the formation of half of 
the amount corresponding to a condition of equilibrium, 80 years 
would be required at a temperature of 725° and one and a fourth 
days at 1225°. At 1825° it takes only 5 seconds. 

These facts are important in determining the heat conditions 
for the preparation of nitric oxide from the air as the first step in 
manufacturing nitrates. Evidently the highest possible tem- 
perature must be secured, and when the nitric oxide has been 
formed at that temperature, it must be cooled as quickly as 
possible through the temperatures at which formation and de- 
composition are both rapid and where the point of equilibrium 
lies farther toward the side of decomposition. 

Nitric oxide is a colorless gas which may be condensed to a 
liquid that boils at —153.6° and freezes at —167°. Is the gas 
heavier or lighter than air? 


1 Nernst, Z. anorg. Chem. 49, 213. 





NITRIC OXIDE 217 


Nitric oxide, quite unlike nitrous oxide, will extinguish a 


-glowing coal or a candle which is thrust into it, or even a piece 


of phosphorus which is barely ignited. If a little hotter, phos- 
phorus burns brilliantly in the gas, and a mixture of the vapor 
of carbon bisulfide with the gas will burn with a bright blue flash, 
which is rich in the light rays that affect a photographic plate. 

Nitric oxide combines directly with oxygen to form nitrous 
anhydride, N2Os, nitrogen dioxide, NOs, or nitrogen tetroxide, 
N.O,, or a mixture of the three, according to the temperature, 
and the relative volumes of the gases. If a strong base, such 
as potassium hydroxide, KOH, is present when nitric oxide and 
oxygen are brought together, a nitrite, KNOs, is formed, even 
though the oxygen is in excess. In what proportion by volume 
must nitric oxide and oxygen be brought together to form nitrous 
anhydride? In what proportion to form nitrogen dioxide? 

The structural formula of nitric oxide is usually written 
N=O, representing the nitrogen as bivalent, a valence which is 
very unusual for the element and which it is not known to have 
in any other compound. This unusual structure seems to be 
closely connected with the tendency of nitric oxide te combine 
with more oxygen. 

* Nitric oxide is formed by the reduction of nitric acid by 
ferrous sulfate, FeSQu, the latter being oxidized to ferric sulfate, 
Fe2(SOx)s, if sulfuric acid is present : 


2 HNO; = 2 NO-+ (3 O) + H:O 
2 FeSO, a H,.SO, -t (O) SS Fe2(SO.)s3 +. H.O 


Combining the equations, 
6 FeSO, of a H.SO, fe Z HNO; Sa Fe(SO.)s3 + 3 H.O + 2 NO 


The nitric oxide is absorbed by a solution of ferrous sulfate 
with the formation of an unstable compound, FeSO..NO 
(see Manchot and Zechentmayer, Ann. 350, 368), which gives 
to the solution a dark brown or black color and which is often 
used for the qualitative detection of nitric acid. The reduction 


218 A TEXTBOOK OF CHEMISTRY 


of nitric acid to nitric oxide by ferrous chloride, FeCle, and hydro- 
chloric acid is also used for its quantitative determination. 

Nitrous anhydride, N2O3:, may be prepared by the union of 
oxygen with nitric oxide, by the action of a dilute acid on sodium 
nitrite, NaNOs, or by the reduction of nitric acid, especially by 
arsenious oxide, As2O3: 


4 NO+ O,2 = 2 N.O3 
hoe NaNO, —+- H.SO, == NaySOx. + y HNO, 
Nitrous 
Acid 
| 2 HNO, = N.O3 + H,O 
Z HNO; + AsoOz3 + 2 H,O el H;AsO, -+ N.O3 


Arsenic 
Acid 


At ordinary temperatures a gas having the composition of 
nitrous anhydride has a density which indicates that the com- 
pound dissociates into nitric oxide, NO, nitrogen dioxide, NOs, 
and nitrogen tetroxide, N2O.: 


N.O; = NO + NOz 
2 NO. == N20, (See below.) 


At a low temperature the gases recombine, in part, and con- 
dense to a dark blue or green liquid, but even at — 90° there is 
still some dissociation (Ramsay). 

Nitrous Acid. If the mixture of gases spoken of in the last 
paragraph is dissolved in cold water, nitrous acid, HNOk, is 
formed : 


NO + NO, + H,0 = 2 HNO, 


The acid is very unstable and can exist only in dilute solutions. 
Salts of nitrous acid are most easily prepared by reducing a ni- 
trate, as, for instance, sodium nitrate, with lead or copper: 


NaNO; + Pb = NaNO, + PbO 
Sodium 
Nitrite 


NITROGEN DIOXIDE 219 


Nitrogen Dioxide, NOs, and Nitrogen Tetroxide, NeO,. 
When nitric oxide is mixed with air or oxygen, the colorless gas 
changes to a reddish brown color and is converted, at ordinary 
temperatures, into a mixture of nitrogen dioxide, NOs, and 
nitrogen tetroxide, N2O4y. The molecular weight of nitrogen 
dioxide is 46, that of nitrogen tetroxide is 92. As the gram 
molecular volume of the gas weighs about 80 grams at 10°, 
69 grams at 64° and 46 grams at 150°, it is evident that at the 
lower temperature it consists chiefly of nitrogen tetroxide, N2Ou, 
but that this dissociates into nitrogen dioxide, NOs, as the 
temperature rises : 


NO, = 2 NO; 


At still higher temperatures nitrogen dioxide dissociates into 
nitric oxide, NO, and oxygen. 


2NO, 2-2 NO + 0, 


The mixture of the two gases may be easily condensed to a 
reddish brown liquid which boils at 25°. At lower temperatures 
the liquid becomes lighter colored and solidifies to colorless crys- 
tals at —10.5°. From this it would seem that nitrogen tetrox- 
ide is colorless, while the dioxide is colored. 

The structure of the two compounds is probably : 


ne ahd _N-0-N=0 


oN 
The tetroxide, DN—-O-N—0, may be considered as 
O 
partly nitrous anhydride, O=N—O—N=O, and partly nitric 
O O 
anhydride, YN—O—NC . In accordance with this view 
oF No 
it gives both nitric and nitrous acid when dissolved in cold 
water : 


220 A TEXTBOOK OF CHEMISTRY 


O 
CA te) = H-O-N¢_+H—O—N=0 
H_O—H No 
l Nitric Acid Nitrous Acid 
In warm water nitrogen dioxide forms nitric acid and nitric 


oxide: Fe 
NO, + HOH = H—O—N + (H) 
No 


2(H)+ NO, = NO+ HO 
or 3 NO, + H,O ==i2 HNO; +- NO 


As the nitric oxide on coming to the air immediately combines 
with oxygen to form the dioxide, these reactions furnish a means 
for converting the nitric oxide obtained from atmospheric nitro- 
gen in the electric arc into nitric acid. 

Nitrogen Pentoxide is formed when nitric acid is mixed with 
phosphorus pentoxide : 

2 HNO; -b P.O; = 2 HPO; ot NO; 
Metaphosphoric 
Acid 

It forms colorless crystals which melt at 29.5°. It decomposes 
easily into nitrogen dioxide and oxygen. With water it gives 
nitric acid. | 

Other Compounds of Nitrogen. The halogens are univalent 
toward hydrogen and form only one compound, each, with that 
element. Hydrochloric acid, HCl, is typical for the group. 

Oxygen, which is bivalent, forms two compounds, with hydro- 
gen H—O—H and H—O—O—H. 


Nitrogen, which is trivalent, forms at least five compounds 


with hydrogen. yal 
Ammonia, NHs, or N—H 
NH 
H H 
Hydrazine,} NeHs, = or Yn 
H H 


1 Nitrogen is called ‘‘azote’’ (deprived of life) in French, and 
many names of nitrogen compounds are derived from the root ‘‘az”’ 
or “azo.” 


HYDROXYLAMINE 221 


GL ierat JN 
Hydronitric acid, NsH, or H—NC¢ | 





Ammonium trinitride, NH, or TESN__NG, | 


H jx N 
Hydrazine trinitride, N;H;, or NN <—_NnC | 
H Ny N 

We shall find later that carbon, which is quadrivalent, forms 
many hundred of compounds with hydrogen. 

Nitrogen also forms a third oxygen acid, hyponitrous acid, 
H.2N.2Os, and a derivative of ammonia, hydroxylamine, NH,OH. 

* Hyponitrous Acid, H.,N.O.. From its formula, nitrous 
oxide, NO, might be looked upon as the anhydride of hyponi- 
trous acid, but it will not combine with water to form the acid 
nor will it combine with bases to form salts of the acid. The 
method of preparation which indicates most clearly the structure 
of the acid is by the interaction of hydroxylamine and nitrous 
acid : he eee 

Pa 
H—O=N + O=N—OH = H—O—N=N—O—H + H.20 
ei a 

The pure acid can be obtained in crystalline form, but is very 
explosive. The lowering of the freezing point of an aqueous 
solution of the acid (see p. 112) shows that the formula is H»N2Oz, 
and not the simpler formula, HNO. In solutions of either 
alkalies or acids it decomposes slowly into water and nitrous 
oxide, N2O. 

* Hydroxylamine, NH.OH, may be prepared by the electro- 
lytic reduction of nitric acid, using an amalgamated lead cathode 
and in the presence of sulfuric acid, which combines with the 
hydroxylamine to form hydroxylammonium sulfate :! 


1 Tafel, Z. anorg. Chem. 31, 289. 


2 
222 A TEXTBOOK OF CHEMISTRY 


O u 
Me ain ie ee 
Heo) NC + oH Hoo NC, + 20 


2 NH.OH + H2SO, = (NH;:0H)2SO, 
Hydroxylammonium 
Sulfate 
Pure hydroxylamine may be obtained in white scales or 
needles which melt at 33°. It dissolves easily in water and the 
solution seems to contain some hydroxylammonium hydroxide, 


N HX , which is a much weaker base than ammonium hy- 
OH 

droxide, NH,OH. Hydroxylamine is much used in organic 

chemistry for the preparation of derivatives of aldehydes and 

ketones, called oximes. 

* Hydrazine, NoH,y. By means of a series of reactions a com- 
pound called bisdiazoacetic acid may be prepared. When this 
is hydrolyzed by warming with hydrochloric acid, it gives hydra- 
zine hydrochloride and oxalic acid. 


aN 
HO,C—CH CHCO.H + 2 HCI + 4 H,0 
Nate 


Bisdiazoacetic Acid 
CO.H 
=2H,N—NH:HCI+2 | 
Hydrazine CO.H 
Hydrochloride Oxalic Acid 


Pure hydrazine is a colorless liquid, which solidifies at a 
low temperature and melts at 1.4°. In solution it forms 
a base, which seems to be H.N—NH;OH, similar to ammo- 
nium hydroxide, NH,OH, and hydroxylammonium hydroxide, 
NH;0H.OH. With acids it forms salts by direct addition, 
taking up either one or two molecules of a monobasic acid, 
the chlorides being H,N—NH3Cl and CIH;N—NH;Cl. The 
second molecule of the acid is held only feebly, however. In 
aqueous solution hydrazine gives a tenfold weaker base than 
ammonium hydroxide. 


4? 


HYDRONITRIC ACID 223 


Hydrazine and its derivatives, especially phenylhydrazine, 
CsH;NHNH2, are much used in organic chemistry for the 
preparation of derivatives of aldehydes, ketones and sugars. 


N 
* Hydronitric Acid or Azoimide, HNC |. By passing am- 
N 


monia over heated metallic sodium a compound called sod- 
amide, NaNHe, may be prepared. When nitrous oxide, N.O, 
is passed over this at a temperature of 190°, sodium trinitride, 


NaN3, is formed: 


iain Hi wit Sie aca N 
Na BS Neale Nolita? 


By dissolving the sodium trinitride in water, adding sulfuric acid, 
and distilling, the hydronitric acid passes over with some water, 
from which it can be separated partly by redistillation and finally 
by treating with calcium chloride: 


2 NaN3 +. H.SO, = NaeSO, + 2, HN3 


The pure acid freezes at — 80 and boils at + 37°. It is very 
explosive, as are also many of its salts, especially the salts with 
silver, AgN3, and copper, Cu(N3)2. The silver and copper salts 
are very difficultly soluble in water. In aqueous solution the free 
acid is only slightly ionized, its strength in this respect resem- 
bling that of acetic acid. It is, however, a much stronger acid 
than hydrogen sulfide, H.S, or carbonic acid, H2CQs. 

Iodine Trinitride. By treating silver trinitride with iodine 
the silver may be replaced by iodine: 


N N 
Ag—NC | +h=Asl + INC \ 


Iodine trinitride is a colorless solid with a penetrating odor, 
resembling that of cyanogen iodide. It is hydrolyzed by water 
to hydronitric and hypoiodous acids : 


224 A TEXTBOOK OF CHEMISTRY 
IN N 
INC | + HOH = H—NC | +1-0-H 
\n N | 


This reaction is remarkable because the iodine conducts itself 
as the positive part of the molecule, combining with the negative 
hydroxyl, while the hydrazoic group, N3, combines with the posi- 
tive hydrogen. Probably for the same reason silver trinitride 
will not react with iodine trinitride to give the molecule Ng, 


n | yn | 
react in the positive form in the two compounds. 

Nitrogen Trichloride and Chloroamine. Chlorine reacts with 
ammonia and ammonium salts forming chloroamine, NHeCl, 
and nitrogen trichloride, NCl3. The reaction, which is rever- 
sible, is most easily explained by the electron theory (p. 181) as 
follows: 


, because both the silver and iodine tend to 





inl H Ht 
Nu: Cl: 2@H:N:Cl:- :Cl:eH: N: Cl: oS H: Cl: 
H H 
Chloroammo- Chloro- 
nium Chloride amine 


It is seen from the above that the molecule of chlorine sepa- 
rates intoa positive chlorine ion, : Ci+ , with six electrons, and a 
negative ion, Cl: :; with eight electron: The former adds 


itself to the nitrogen atom giving a chloroammonium ion, with 
which the negative chloride ion remains momentarily in com- 
bination. It immediately separates, however, with a hydrogen 
atom, as hydrochloric acid. If the solution contains little free 
hydrochloric acid, a second molecule of chlorine reacts with the 
chloroamine in a similar manner, giving dichloroamine, NHC}. 
A third addition of chlorine and loss of hydrochloric acid gives 
nitrogen trichloride, NCl3. 

In the presence of concentrated hydrochloric acid, the reverse 


ENDOTHERMIC COMPOUNDS 229 


reactions occur. Nitrogen trichloride and hydrochloric acid 
give, quantitatively, as the final products, ammonium chloride 
and free chlorine. 

The formation and decomposition of nitrogen trichloride show 
that its chlorine has a very strong tendency to react in the 
positive form. Because of this, one molecule of nitrogen tri- 
chloride is equivalent to six atoms (or three molecules, 3 Cl.) of 
available chlorine in oxidizing power: 


ee NCI; + 3 As9O3 + ks, H.O — 6 H;AsQO, -f- 2 NH; “be 6 HCl 
6 Cl. -f- 3 AseO3 + 15 HO = 6 H3AsO, + ibs HCl 


Nitrogen chloride is a volatile oil, which is very explosive and 
dangerous. Dulong, who discovered it, lost an eye and three 
fingers while working with it, and both Faraday and Davy were | 
injured while experimenting with the substance. It is soluble 
in benzene and may be handled more safely in such a solution. 

* Nitrogen Iodide, NoH3lIs, is formed when a strong solution 
of ammonia is poured over powdered iodine. When dried in the 
air it forms a black powder which explodes violently at the 
lightest touch. 

Endothermic Compounds. ‘The heat of formation of nitrogen 
chloride from its elements is — 54,700 calories (Noyes and 
Tuley). The explosive character of nitrogen chloride and nitro- 
gen iodide is evidently closely connected with the fact that they 
are endothermic compounds, that is, compounds which are formed 
from their elements with absorption of heat and which, con- 
versely, decompose into their elements with evolution of heat. 
The absorption of heat in their formation indicates that the 
atoms have little affinity for each other in the compounds and 
so may easily separate, and the heat generated by the decompo- 
sition, when it once begins, raises the mixture rapidly to a higher 
temperature and this in turn hastens the reaction till it becomes 
explosive. The explosion is further caused, of course, by the 
formation of a large volume of gas at a high temperature from 
a small volume of a liquid or solid. 


226 A TEXTBOOK OF CHEMISTRY 


EXERCISES 


1. The specific gravity of mercury at 0° is 13.6. What is the weight 
of air above one square centimeter of the earth’s surface at sea level ? 

2. Sketch an apparatus suitable for the preparation of nitrogen by 
the use of hydrogen and copper oxide. 


3. What is the weight of a gram molecular volume of air if it contains 
21 per cent of oxygen, 0.9 per cent of argon and 78.1 per cent of nitro- 
gen? What is the weight of one liter ? 


4, If a gram molecular volume of the mixture of nitrogen peroxide 
and nitrogen tetroxide weighs 69 grams at 64°, what is the per cent of 
each compound present ? 


5. How many grams of nitric acid will be required to dissolve 5 
grams of copper? How many liters of nitric oxide will be formed ? 


6. What weight of ammonia will be contained in one liter of aqua 
ammonia having a specific gravity of 0.90 and containing 28 per cent of 
ammonia? How much water? How many grams of ammonia would 
a liter of water take up in forming a solution of specific gravity 0.90? 
How many liters of the gas? 

7. How much salt, NaCl, will be required to furnish the hydrochloric 
acid necessary to neutralize 17 grams of ammonia? How much sul- 
furic acid ? 

8. If an isomeric nitrogen trichloride, NCls, could be prepared in 
which the chlorine atoms were negative, what would be the products 
of its hydrolysis ? 


CHAPTER XIII 


THE ATMOSPHERE. NOBLE GASES 


Determination of Oxgyen. The first determination of the 
per cent of oxygen in the atmosphere was made by Lavoisier, 
who heated a measured quantity of air with mercury as long as 


the oxygen continued to be absorbed (p. 19). 
He determined roughly the contraction which 
occurred and also showed that by heating the 
oxide of mercury formed he obtained a volume 
of oxygen closely agreeing with this contrac- 
tion. His results showed that approximately 
one fifth of the volume of air is oxygen. The 
determination may be made more accurately 
by measuring a volume of air in a eudiometer, 
exposing it to the action of a stick of phos- 
phorus, which is slowly oxidized to phospho- 
rous acid, H3PO3, and measuring the volume 
of the gases which remain (Fig. 75). Or a 
measured volume of air may be mixed with a 
little more than two fifths of its volume of 
hydrogen and the mixture exploded. One 
third of the contraction in volume (or, more 





accurately , see p. 68) will be the volume 


1 
3.0021 
of the oxygen in the air taken. 

Composition of Air. Very many careful 
analyses of air have shown that when samples 
are taken out of doors under usual conditions 









Paonia afas Fan pay 
muinal 


TRRRROOION 


Fig. 75 


the composition of dry air agrees very closely with the fol- 


lowing : 
227 


228 A TEXTBOOK OF CHEMISTRY 











By VoLUME By WEIGHT 
COXV On Oytune. seat aio ea pete 20.95 per cent 23.15 per cent 
Carbon cioxides. "00. wae 0.03 per cent 0.05 per cent 
AT QOR. co eee sates eta eee 0.94 per cent 1.3 per cent 
NItrOperl nee ae ree ee 78.08 per cent 75.5 per cent 
100 100 








The percentage of oxygen varies slightly, but is very rarely 
less than 20.9 or more than 21.0 per cent, out of doors. 

Air is a Mixture. Although the composition of air is nearly 
constant, it is believed that the oxygen and nitrogen which it 
contains are merely mixed together and not chemically com- 
bined, for the following reasons : 

1. No heat is generated when oxygen, nitrogen, argon and 
carbon dioxide are mixed, and yet the mixture has all of the 
properties of air. 

2. Liquefied air when it boils gives at first a gas richer in 
nitrogen and at last nearly pure oxygen. A compound when 
it boils away without decomposition has the same composition 
from first to last. i 

3. When water is shaken with air, it dissolves the oxygen and 
nitrogen in proportion to the solubility and partial pressure 
of each, and as oxygen is more soluble than nitrogen, the mixture 
of gases obtained by boiling the water contains about 35 per cent 
of oxygen, while air contains only 21 per cent. This is the 
conduct to be expected of a mixture rather than that of a com- 
pound. 

4. A gram molecular volume of air weighs 28.95 grams. The 
simplest formula which would give a composition approximating 
that of air is N4O. The gram molecular volume of a compound 
of this formula would weigh 72 grams. 

On the other hand, a liter of air weighs 1.2928 grams, which is 
almost exactly what it should weigh if it is a mixture of gases 
in the proportion which has been given, as will be seen from the 
following calculation : 


THE ATMOSPHERE 229 


WEIGHT OF ONE 


LITER 
Oxygen 1.429 X 0.2095 = 0.2994 
Carbon dioxide 1.9768 X 0.0003 = 0.0006 
Argon 1.7828 X 0.0094 = 0.0168 
Nitrogen 1.2507 X 0.7808 = 0.9765 


1.2933 


Carbon Dioxide in the Air. The carbon dioxide in the air 
comes from four principal sources: 1. From the breath of men 
and animals, the carbon of food which is eaten being mostly 
converted into carbon dioxide by the oxidation processes which 
take place in the body. 2. From the burning of compounds 
containing carbon, such as wood, coal, oil or natural gas. 
3. From the decay of animal and vegetable substances under 
the influence of bacteria. 4. From volcanoes and other sub- 
terranean sources. 

While the amount of carbon dioxide in the air is very small 
in comparison with the amounts of oxygen and nitrogen, the total 
amount is very great. Thus it is estimated that 1,300,000,000 
tons of coal are burned annually, and this gives nearly three 
times its weight of carbon dioxide, but this would increase the 
amount in the air by only one six-hundredth part.1. The carbon 
dioxide of the air is also a very important factor in the economy 
of nature, as it furnishes practically all of the carbon for the 
growth of plants. In a sense it is the constituent of the atmos- 
phere which is most vitally important for the life of plants, 
as oxygen is the constituent necessary for the life of animals. 
In another sense it may be said to be even more important, since 
it furnishes the most important constituent for the growth of 
plants, while oxygen furnishes to animals only the means with 
which to consume food which is secured from some other source. 
In utilizing carbon dioxide plants decompose it and exhale oxygen 
to the air. In this way the growth of plants prevents the accu- 
mulation of carbon dioxide in the atmosphere. The process of 


1A. Krogh, quoted by F. W. Clarke in Data of Geochemistry, 
p. 42. See also Science, 1911, p. 757. 


230 A TEXTBOOK OF CHEMISTRY 


reducing carbon dioxide to the compounds of carbon synthe- 
sized by the plant is endothermic, of course, and the necessary. 
energy is furnished by the sunlight. When we burn coal in our 
furnaces, we make use of the energy of sunlight which was stored 
by growing plants millions of years ago. It has even been sug- 
gested that all of the oxygen of the air came originally from car- 
bon dioxide through this process of plant growth (Lord Kelvin). 

Besides the equilibrium maintained by the balance between 
the evolution of carbon dioxide from the sources named and the 
absorption of the gas by growing plants, the ocean plays a very 
important part in maintaining a constant amount in the air 
during long periods. As carbon dioxide dissolves in water in 
proportion to the partial pressure of the gas (p. 165), any increase 
in the amount of carbon dioxide in the air would be followed 
very quickly by an increase in the amount in the ocean, while 
any decrease would be replaced from the storehouse in the ocean. 
Since the amount of “ free’”’ carbon dioxide in the ocean to a 
depth of 5 kilometers is nearly fifty times ' the amount in the air 
above it, and about three fourths of the earth’s surface is covered 
by the ocean, the importance of the store contained in the ocean 
is obvious. 

Ventilation. Shortly after the discovery of the composition 
of the air by Priestly and Lavoisier, a method was devised for 
analyzing air by mixing it with nitric oxide, to combine with the 
oxygen, and then absorbing the nitrogen peroxide formed by 
means of a solution of potassium hydroxide. It will be readily 
understood that such a method requires very great care to secure 
accurate results and the early determinations led the observers 
to think that there was a considerable fluctuation in the amount 
of oxygen present and that this fluctuation caused the difference 
between good and bad air. But Cavendish was able to use even 
the nitric oxide method so accurately that he very soon showed 
that the variation in the composition of the atmosphere must 
be between very narrow limits, and this result has been confirmed 


1 Calculated on the basis of 45 milligrams per liter of sea water. 


VENTILATION Bat 


by later observers. It was then discovered that carbon dioxide 
in mines and in wells or caves frequently killed persons exposed 
to its action, and for many years it was supposed that this gas 
acts as a positive poison and is the chief cause of danger in poorly 
ventilated rooms. This fallacy and also the opinion that carbon 
dioxide will accumulate near the floor of a room because the gas 
is one and a half times as heavy as air were spread so widely in 
semipopular literature and became so firmly fixed in the minds 
of many people that it has proved very difficult to correct these 
errors. It has been shown that the amount of carbon dioxide 
present in the air of even badly ventilated rooms is practically 
never great enough to cause any injury to human beings. It has 
been found very difficult to demonstrate clearly just what sub- 
stances cause the ill effects which follow from poor ventilation, 
and some recent authorities have spoken doubtfully of the 
standards for ventilation which have been proposed. There 
seems to be little doubt, however, that lack of ventilation in 
factories, offices and dwellings is a frequent cause of disease. It 
is also very well established that abundance of fresh air secured 
by life out of doors, both by night and day, combined with a 
nourishing diet, furnish the best hope of recovery from incipient 
tuberculosis. 

While exhaled carbon dioxide is not in itself harmful, it fur- 
nishes the best means of determining whether a room occupied 
by people is properly ventilated or not. The amount of the 
gas should not exceed 0.07 per cent by volume. To secure 
this amount of ventilation 55,000 liters or 2000 cubic feet of 
fresh air will be required each hour for each person in a room.} 

Moisture. Natural air always contains a certain amount of 
water vapor, but this is subject to very great variations, depen- 
dent on the temperature and the conditions to which the air has 
been subjected. The pressure of the water vapor can never 
much exceed the normal vapor pressure of water for the given 


1 Roscoe and Schorlemmer, Treatise on Chemistry, I, 589. 
Another authority recommends 85,000 liters per hour. Stewart, 
Manual of Physiology, p. 244. 


232 A TEXTBOOK OF CHEMISTRY 


temperature, and, indeed, can only exceed that when in a state 
of unstable equilibrium such that the introduction of suitable 
nuclei to form points of condensation will at once cause the for- 
mation of a cloud. When the pressure of the water vapor in the air 
corresponds to the normal pressure of water vapor for the given 
temperature, the air is said to be saturated, but such a condition 
does not usually obtain close to the earth’s surface. At a height 
of a few hundred or thousand feet, however, owing partly to the 
mixing of warm, nearly saturated air with colder air currents, 
partly to the lowering of the temperature, which results from 
the adiabatic cooling of air as it expands in rising, saturation 
and condensation to clouds and rain take place. 

The amount of moisture in air may be determined: 1. By 
aspirating a known volume through weighed bulbs containing 
concentrated sulfuric acid. 2. By determining the dew point, 
that is, the temperature at which the air will deposit moisture on 
a cooled, polished, metallic surface or the temperature at which 
moisture will just disappear from such a surface. 3. By com- 
paring the temperature of the air with the temperature of a 
thermometer whose bulb is covered with moist cotton over which 
air is blown. ‘Tables have been prepared giving the humidity 
corresponding to the difference observed. 

In general the humidity in rooms which are heated is too low 
for healthfulness and should be supplemented by artificial means. 
In many factories, especially in those for spinning and weaving, 
the degree of humidity is of vital importance to the success of the 
operations. 

Liquid Air. Critical Temperature. After it had been shown 
early in the nineteenth century that such gases as chlorine, 
ammonia, carbon dioxide, and sulfur dioxide could be liquefied, 
many attempts were made to liquefy oxygen and nitrogen, or 
air. These gases were subjected to pressures of several hundreds 
of atmospheres, but it was always found that the gas continued 
to fill completely and uniformly any space left to it, while if 
liquefied it should have separated into a liquid and a gaseous 
portion. Finally, in 1869, Andrews showed that carbon dioxide, 


CRITICAL TEMPERATURE 200 


which can be liquefied under a pressure of 38.5 atmospheres at 
0°, or 71 atmospheres at 30°, cannot be liquefied, even under 
pressures very much greater than this, at temperatures above 31°. 
If a thick-walled, sealed glass tube containing liquid carbon diox- 
ide is warmed gently, at a temperature of 31.35° the liquid in the 
lower part of the tube will suddenly disappear and the gas will 
now fill the tube uniformly. The pressure may be increased or 
decreased, but as long as the temperature is above 31.35° no pres- 
sure either high or low can be found at which the carbon dioxide 
will separate into a liquid phase and a vapor phase. Below this 
temperature carbon dioxide will be partly liquid and partly gas, 
provided the pressure is equal to the vapor pressure of the liquid 
at the given temperature and the volume filled by the substance 
is large enough to allow a part to assume the vapor phase. The 
temperature above which a gas cannot be liquefied is called the 
critical temperature. 

Andrews’ experiment made it seem very probable that the fail- 
ures to liquefy air were due to the fact that the critical tempera- 
tures of oxygen and nitrogen are much below ordinary tempera- 
ture. Following this suggestion, Cailletet in Paris and Pictet 
in Geneva (1877), working independently, both succeeded in 
liquefying oxygen by the use of cold and pressure combined with 
the cooling effect produced by the expansion of the highly com- 
pressed gas. 

Some years later it was shown by Joule and Thomson (Lord 
Kelvin) that a moderately compressed gas scarcely changes its 
temperature on expanding into a vacuum — for instance, if air 
compressed to 20 atmospheres is allowed to expand into a vacu- 
ous receptacle, both receptacles being surrounded by water, the 
temperature scarcely changes, though for all gases except hydro- 
gen and helium, there is a slight cooling effect. This cooling effect 
increases for higher pressures, or when a gas is so far compressed 
that it no longer obeys Boyle’s law (p. 35). It would seem that 
the attraction between the molecules of the gas has a greater 
effect as the molecules are brought closer together, causing the 
gas to contract more than it should in accordance with the law. 


234 A TEXTBOOK OF CHEMISTRY 


When the gas expands from such a condition work must be done 
in overcoming this attraction between the molecules, and the 
expansion is accompanied by a cooling effect. On the basis of 
these facts Linde, Hampson and others have devised machines 
by means of which air can be readily liquefied in large quantities. 
In these machines air is compressed to 
150-200 atmospheres and is then allowed 
to expand to atmospheric pressure in such 
a manner that the expanded and cooled 
air passes back over the tube in which 
the air is expanding. In the Hampson 
machine the air expands through a copper 
tube of about three millimeters in internal 
diameter and one hundred and _ thirty 
meters in length. This is wound in a 
spiral to secure compactness and the ex- 
panded air is compelled to follow the 
course of the spiral backwards, Fig. 76. 
By these machines a portion of the air is 
soon cooled to the point of liquefaction 
and the liquid air collects in a receptacle 
placed beneath the end of the spiral. 

The carbon dioxide must be removed 
from the air which is to be liquefied, by 
passing it through a large apparatus filled 
with slaked lime, and the moisture must 
also be removed by calcium chloride or 
some drying agent, as otherwise these 
would condense in solid form and stop up 
the tube through which the air expands. For the liquefaction of 
hydrogen the compressed gas must be cooled by liquid air, as 
it is only at low temperatures that hydrogen departs sufficiently 
from Boyle’s law so that it can be liquefied by this method. 

Liquid nitrogen boils at —194°, liquid oxygen at —182.5°. 
Liquid air will contain, therefore, a larger proportion of oxygen 
than ordinary air, and by a sort of fractional distillation it is easy 

















LIQUID AIR. ARGON 230 


to obtain from it a gas which contains from 75 to 95 per cent of 
oxygen. Such a gas is used for medicinal (e.g. in pneumonia) 
and some technical purposes. The method 

is also used to obtain nearly pure nitrogen 

and is now the most important industrial 

method for the preparation of both oxygen 

and nitrogen. 

For experimental purposes liquid air is 
kept in Dewar flasks (Fig. 77), double- 
walled flasks having the space between the 
two walls evacuated to prevent loss of heat 
by convection currents. The inner bulb is 
often silvered to cause it to reflect radiant Fig. 77 
heat which reaches it from outside. 

Argon, A, 39.91. In 1785 Cavendish described an experiment 
in which he mixed air with an excess of oxygen, passed electric 
sparks through the mixture, and absorbed the oxides of nitrogen 
formed by a solution of potassium hydroxide. He then absorbed 
the rest of the oxygen by means of “liver of sulfur” and re- 
ported that the gas remaining unabsorbed was not more than 
+3 of the original volume of the air. The real significance of 
this remarkable experiment was not understood for more than 
a century. 

During the eighties and nineties of the last century Lord 
Rayleigh spent a great deal of time in determining very accu- 
rately the density of the elementary gases, oxygen, hydrogen 
and nitrogen. In the course of his work he prepared what he sup- 
posed to be nitrogen by removing oxygen and all other known 
substances from the air. He also prepared nitrogen by the 
decomposition of ammonia. To his surprise a liter of the nitro- 
gen obtained from the air weighed about 6 milligrams more than 
a liter of nitrogen prepared from ammonia. Lord Rayleigh is 
a physicist, and he called in the assistance of a chemist, Sir 
William Ramsay, to solve the problem which was _ presented. 
Within a short time, in 1894, the two succeeded in preparing 
argon, partly by a repetition of the Cavendish experiment with 


236 A TEXTBOOK OF CHEMISTRY 


modern appliances, partly by removing the nitrogen of the air 
by passing it over heated magnesium, with which the nitrogen 
combined. 

Argon was not only a new element, but it belongs to a wholly 
new class of elements, now called the Zero group of the Periodic 
System, or the noble gases. The most remarkable property of 
these elements is that none of them enters into chemical combina- 
tion with other elements or with itself — the valence of the 
group is zero. 

Argon may be condensed to a liquid, which freezes at — 188° 
and boils at —186.1°.. The gram molecular volume weighs 
39.9 grams. From this the molecular weight is 39.88. 

Atomic Weight of Argon. Specific Heat of Gases. The fact 
that argon will not combine with any other element would, of 
itself, lead us to expect that the molecule of argon consists of a 
single atom and that the formula of the gas is A and the atomic 
weight 39.88. Another, wholly independent, line of evidence 
points to the same conclusion. The specific heat of a gas may be 
determined either while the volume of the gas remains constant 
or while the pressure remains constant. It is evident that the 
specific heat must be greater at constant pressure than at con- 
stant volume because at constant pressure the gas must expand 
as it grows warm and do work as it expands against the pressure 
of the atmosphere. It can be shown that, on the basis of the 
fundamental assumptions of the kinetic theory of gases, in any 
gas in which the energy required to increase the temperature 
of the gas is all used in increasing the average velocity of the mole- 
cules, the ratio of the specific heats must be: 


Specific heat at constant pressure — 1.67 


Specific heat at constant volume 1 
On the other hand, if a part of the energy is used in causing 
the atoms within these molecules to vibrate more violently, 
the numerator of the fraction expressing the ratio between the 
two kinds of specific heat will be smaller, since both kinds of 
specific heat will be greater and an addition to both the numera- 


HELIUM Zab 


tor and denominator of any fraction causes it to approach 
unity. 

The ratio between the two specific heats can be calculated from 
the velocity of sound in the gas. The two kinds of specific heat 
have also been determined directly for air and some other gases. 
It has been found that the ratio of the specific heats for mercury 
vapor, for argon and for some other gases is very close to 1.67/1 
and it is believed that all of these gases are monatomic. The 
ratio of the specific heats for nitrogen is 1.41/1; for carbon 
dioxide it is 1.305/1; for ethylene, 1.26/1; and in general the 
numerator becomes smaller as the molecule is more complex. 
This seems to mean that in gases with complex molecules a con- 
siderable part of the energy used in heating the gas is absorbed 
in doing internal work in the molecules, that is, in causing their 
atoms to vibrate more and more rapidly. It will be readily seen 
that this conclusion gives a simple explanation of the fact that 
complex molecules are generally unstable at high temperatures. 

Helium, He, 4.00. In 1868 Lockyer observed some bright 
lines in the spectrum of the corona of the sun, which did not 
correspond to the lines of any element then known. He called 
the element which gives these lines helium (from #Avos, the sun), 
and he had, in reality, discovered a new element, which for 
nearly thirty years was known to exist only in the sun, 90,000,- 
000 miles away. Shortly after the discovery of argon it was 
recalled that Dr. Hillebrand of the U. S. Geological Survey had 
obtained a gas from the mineral uraninite. Ramsay, on further 
examination of the gases obtained from cleveite, a variety of 
uraninite, found in them not only a small amount of argon, 
but also a gas which gave the same spectral lines which had 
been observed in the light of the sun’s corona, and he soon 
separated helium from the mixture. Helium is only twice as 
heavy as hydrogen and has the lowest boiling point of any known | 
substance (unless we call the electron an element). It boils 
at —268.5° or at 4.5° absolute. Helium has acquired a very 
extraordinary interest, also, from the discovery that it is formed 
by the decomposition of radium. In spite of this method of 


238 A TEXTBOOK OF CHEMISTRY 


formation, radium cannot be considered as a compound of helium, 
and no one has been able to induce helium to combine with any 
other element. 

Helium is found in all gases issuing from the earth. It is 
doubtless derived from radium and other radioactive elements. 
It has been suggested that the reason why only a very minute 
quantity of helium is found in the atmosphere is because, owing 
to the lightness of the helium atoms, their kinetic velocity is such 
that they may fly away from the earth into space. 

Neon, Krypton, Xenon, and Radon. The following partial 
table of atomic weights taken from the periodic system indicates 
that there should be three or four other elements belonging to 
the same family as helium and argon: 


ee He 4 Li i 
FB.) 19 — Na 23 
Cl ae3).3 A 39.9 Gr pt!) 
Br 80 —— Rb_ 85.4 
Le fal 2% a Gs 41335 


A systematic search for these elements soon led Ramsay to the 
discovery of neon (Ne = 20.2), krypton (Kr = 82.92) and 
xenon (X = 130.2) as constituents of the air, each of them pres- 
ent, however, in only very small amounts. Several years later 
it was shown that an evanescent element formed by the disinte- 
gration of radium belongs to this series. The density of the gas 
has been determined only approximately because of the minute 
quantity which it is possible to obtain. From this determina- 
tion the atomic weight is about 222. The element is called 
radon. It disintegrates spontaneously and very rapidly, one 
half of it disappearing in a little less than four days. 

The existence of radon was first established by Dorn in 
1900. It was called radium emanation till Ramsay and his 
co-workers established its physical properties. See Cameron, 
Radiochemistry, p. 47. 


a — 


THE ATMOSPHERE. NOBLE GASES 239 


EXERCISES 


1. An adult eats food containing about 300 grams of carbon daily. 
If this is exhaled as carbon dioxide, CO», at a temperature of 37°, how 
many liters of the gas are exhaled per hour? 


2. If a person breathes 20 times per minute, 500 cc. of air being ex- 
haled at each respiration and the exhaled air contains 4 per cent of 
carbon dioxide, how many liters of the gas are exhaled per hour? 


3. How does the volume of carbon dioxide from a gas jet burning 
3 cubic feet of gas per hour and giving an equal volume of carbon dioxide 
compare with that exhaled by an adult? 


4. If outside air contains 0.03 per cent of carbon dioxide, how often 
must the air in a room 5 meters square and 3 meters high be changed 
in order that the amount of carbon dioxide may not exceed 0.07 per cent 
when two persons, each breathing out carbon dioxide at the rate of 20 
liters per hour, are present ? 


CHAPTER XIV 
PHOSPHORUS 


THE atomic weights of the nonmetallic elements of the fifth, 
sixth, seventh and zero groups of the periodic system and of the 
semimetallic elements of the fifth group are, in round numbers : 




















FirtH Group SrxtH Group SEVENTH GROUP ZERO GROUP 
He4 
N 14 216 F 19 Ne 20 
Pral 5 32 Clo. A 40 
As 75 Se 78 Br 80 Kr 83 
Sb 120 Te 127.5 a ere Xe 130 
Bi 208 woe Nt 222 











Phosphorus, P, 31.03. Occurrence. Phosphorus, the second 
element of the fifth group, is a very important element both for 
vegetable and animal life. It is an essential mineral constituent 
in soils for the growth of plants, and it is also an important ele- 
ment in the protoplasm of the cells and in the bones of animals. 
The ash left when the organic matter is burned out of bones con- 
sists very largely of calcium phosphate, Ca3(PO.)2. The same 
compound is found mixed with other substances in extensive de- 
posits of “phosphate rock” in North and South Carolina, 
Georgia, Florida and Tennessee. These deposits are extensively 
mined for use in applying to soils which are deficient in phos- 
phorus. Phosphorus is also found in the mineral apatite, 
Cas(POx.)3F or Cas(POx)3Cl, which has already been mentioned 
in connection with fluorine. Phosphorus compounds are found 


in almost all iron ores, lessening their value when present in 
240 


‘. 


PHOSPHORUS 241 


more than very small amounts, because of the injurious effect of 
the phosphorus on the iron made from such ores. 

Preparation of Phosphorus. When a. mixture of sand (sili- 
con dioxide, SiOz), calcium phosphate, Ca3(PO.)s, and charcoal 
or coke, C, is heated to a very high temperature in an electric 
furnace, calcium silicate, CaSiO3, phosphorus, P4, and carbon 
monoxide, CO, are produced : 


Ca3(PO,)> a 3 SiO» +. 5 C = 3 CaSiO3 oh De Put 5 CO . 


The phosphorus distills from the retort in which the mixture 
is heated and is condensed and collected under water. This 
electrical furnace method for manufacturing phosphorus has dis- 
placed older, more complicated methods, in comparatively 
recent times. 

Allotropic Forms of Phosphorus. The phosphorus obtained as 
described is a waxlike solid which usually has a slight yellow 
color and this form is called “ ordinary ” or “ yellow’ phos- 
phorus. When pure it melts at 44.5° and can be readily melted 
and cast into sticks under water. Its specific gravity is 1.8232 
at 20°. It boils at 290°. In the gaseous form a gram molecular 
volume weighs about 124 grams, from which the formula must 
be j Pers 

Ordinary phosphorus glows with a pale light when exposed to 
moist air. It may be distilled with steam, and a very minute 
quantity may be detected in a dark room by the use of these 
properties. The word phosphorescence recalls, of course, the 
luminous quality of the element. 

If ordinary phosphorus is heated to 240°-250° in a closed 
vessel, it is gradually, though not quite completely, transformed 
into the allotropic variety called red phosphorus. This was 
formerly called amorphous phosphorus, but it may be crystallized 
from solution in melted lead. When pure and free from yellow 

1 In writing this equation notice that one molecule of calcium 
phosphate, Cas(PO.)2, gives two atoms of phosphorus. The rest of 


the equation follows logically from the formulas of the products 
formed. 


242 A TEXTBOOK OF CHEMISTRY 


phosphorus its specific gravity is 2.34. It is not poisonous, 
while yellow phosphorous is very poisonous indeed. Yellow 
phosphorus dissolves readily in carbon disulfide, red phosphorus 
does not. Yellow phosphorus must be kept away from the air 
and is usually kept under water because of the very low kindling 
temperature. Red phosphorus takes fire at a much higher tem- 
perature and may be kept in open bottles. When heated to a 
high temperature, red phosphorus distills and goes back to the 
yellow form, but at lower temperatures the vapor pressure of 
red phosphorus is much lower than that of the yellow variety. 
The molecular weight of red phosphorus has not been deter- 
mined. 

Matches. The methods of obtaining fire in use before the nine- 
teenth century were difficult of application and people often sent 
to their neighbors even at some distance for coals rather than to 
take the trouble of starting a new fire. Phosphorus was dis- 
covered, it is true, in 1669 by Brandt, an alchemist of Hamburg, 
but it was not till 1827 that use was made of its low kindling 
temperature for the preparation of matches. For the first 
matches using phosphorus the match sticks were dipped in 
melted sulfur and then in a mixture of phosphorus and glue 
or some other adhesive substance. When dry a slight friction 
raises the phosphorus to its kindling temperature and this, as it 
burns, sets fire to the sulfur, which, in turn, ignites the wood. 
In the later manufacture the sulfur was replaced by other com- 
bustible substances which do not give an objectionable odor, and 
the kindling power of the phosphorus was reenforced by potas- 
sium chlorate, red lead or other oxidizing compounds. Ordinary 
phosphorus is extremely poisonous, however, and gives off enough 
vapor at ordinary temperatures so that, unless extraordinary 
pains are taken to ventilate the factories, the workmen often 
suffer from a very painful and fatal disease, which causes necro- 
sis of the jaw. Partly for this reason and partly to avoid the 
danger of accidental fires, most European countries have for- 
bidden the sale or even the manufacture of matches containing 
ordinary phosphorus. The “ safety’ matches used in these 


PHOSPHINE 243 


countries have on their heads, usually, a mixture of antimony 
trisulfide, potassium chlorate and glue, and they are ignited 
on a prepared surface of red phosphorus, glue and a sulfide of 
antimony. 

In comparatively recent times it has been discovered that 
tetraphosphorus trisulfide, P453, may be substituted for yellow 
phosphorus in ordinary matches. As it does not give off poison- 
ous vapors, this sulfide of phosphorus ought soon to entirely 
displace the ordinary phosphorus for this manufacture. A law 
passed by Congress in 1912 will prevent the further use of ordi- 
nary phosphorus for matches in the United States. 

Phosphine, PH;. When yellow phosphorus is warmed with 
a strong solution of sodium hydroxide, it is oxidized to sodium 
hypophosphite, NaH,PO.. At the same time some of the hy- 
drogen of the water or of the sodium hydroxide combines with 
more of the phosphorus to form phosphine, PH;. In preparing 
the gas a rather small flask should be used, and it is well to add 
to the contents of the flask, before warming, a few drops of 
ether, which will expel the air and prevent a possible explosion. 
The phosphine prepared in this manner contains some hydrogen 
and some of the liquid hydrogen phosphide, P2H4, which corres- 
ponds in composition to hydrazine, NeH4. This liquid is volatile 
and takes fire spontaneously on exposure to the air. For this 
reason, although the kindling temperature of phosphine is about 
150°, the phosphine prepared as described takes fire at once as 
it comes to the air. Bubbles of the gas explode as they come to 
the surface of the water, forming a cloud of phosphoric acid, 
H3PO,, which gives beautiful vortex rings in still air (Fig. 78). 
Phosphine may be condensed to a colorless liquid, which boils 
at — 86.2° and solidifies at lower temperatures to crystals which 
melt at — 133°. 

Phosphonium Salts. Phosphine combines with acids to form 
phosphonium salts, as ammonia forms ammonium salts. The 
most stable and best known of these salts is phosphonium iodide, 
PHI, which may be prepared by the direct union of phosphine, 
PH3, and hydriodic acid, HI. It forms white crystals, which 


244 A TEXTBOOK OF CHEMISTRY 


sublime at 80°. The salt is hydrolyzed by water and phosphine 
escapes from the solution: 


PH,I -+ HOH = PH,OH + HI 
PH,OH = PH; + HOH 


Evidently the phosphonium group, PHg, is very unstable, even 
in the presence of hydrogen ions. It is sometimes stated that 
phosphine is a much weaker base than ammonia. Correctly 
speaking neither is 
a base, and the true 
base, phosphonium 
hydroxide, PH,OH, 
is extremely un- 
stable, if it exists 
at all. We shall 
find that arsine, 
AsH3, and stibine, 
SbH3, do not com- 
bine with acids. 
In the series NHs, 
PH, AsHs, SbHs3, 
not only does the 
tendency to com- 
bine with a fourth 
hydrogen atom be- 
come less and less, but the compounds themselves are less and 
less stable, stibine, SbH3, decomposing at ordinary temperatures, 
especially in the presence of metallic antimony. 

Phosphorus Trichloride, PCl3, and Phosphorus Pentachloride, 
PCl;, are easily prepared by the direct union of chlorine and 
phosphorus. The trichloride is a liquid which boils at 76°. 
The pentachloride is a white solid which melts in a sealed tube 
at 148°. Its vapor pressure, however, is 760 mm. at 140°. In 
other words its melting point is higher than its boiling point and 
it sublimes without melting when heated under atmospheric 
pressure. 





a2 


CHLORIDES OF PHOSPHORUS 245 


The molecular weight of phosphorus pentachloride, PCl,, 
is 208.5; but a gram molecular volume of the gas at 182° weighs 
147 grams, while at 300° it weighs only 105.7 grams, only a little 
more than one half the weight of a gram molecule. This indi- 
cates that the pentachloride dissociates into phosphorus trichlo- 
ride and chlorine, the dissociation being nearly complete at 300°. 
This gives twice as many molecules as there are in the original 
pentachloride and one gram molecule of the pentachloride gives 
two gram molecular volumes of gas: 


PCL PC - Cl 


Hydrolysis of the Chlorides of Phosphorus. The chlorides 
of phosphorus are decomposed, or hydrolyzed by water in the 
same manner as most other chlorides of nonmetallic elements : 

PCl; + 3 H.OH = H;PO; + 3 HCl 


Phosphorous 
Acid 


PC], + 4H.OH = H3PO, + 5 HCl 
Phosphoric 


Acid 
Phosphorus Oxychloride, POCl;. When phosphorus penta- 
chloride is treated with a small amount of water or with almost 
any compound containing the hydroxyl group, OH, it is changed 
to phosphorus oxychloride, POCI;, while the two chlorine atoms 
which are lost combine with the two atoms which were united 
to the oxygen: 


Cl Cl Cl 
Pol Lads) ae o-Pocl es a 
Cl Cl Cas 
Cl Cl Wel 
yPccl + CHOH = O=POCl+ pe ng 
Cl Cl nsayiitesdtetcl ie) Cl pinyl Chloride 
Cl Cl Cl 
| Dame 
YpLal + GMO = O=PLC+ Croc 
Cl Cl (or C,H;0.0H) Cle set Chlonde 


Acetic Acid 


246 A TEXTBOOK OF CHEMISTRY 


Phosphorus oxychloride may also be prepared by oxidizing 
phosphorus trichloride with potassium chlorate. It is a color- 
less liquid which boils at 107.2°. It has a very unpleasant odor 
and the vapor attacks the eyes strongly. It is, of course, hy- 
drolyzed by water to phosphoric and hydrochloric acids. It is 
the chloride of phosphoric acid, H3PO,, in the same sense that 
sulfuryl chloride, SOsCl, is the chloride of sulfuric acid, H2S0, 
(p. 189). 

Oxides of Phosphorus. When phosphorus is burned with an 
insufficient supply of air, a mixture of two oxides, phosphorus 
“trioxide,” P4O¢, and phosphorus “ pentoxide,’ P4O,, 1s formed. 
The names were given long before determinations of the density 
of the vapors of these compounds showed that they have the 
formulas given. The names refer, of course, to the simple formu- 
las P.O; and P2,O;. The trioxide, PsOg, is a solid which melts at 
22.5° and boils at 173.1°. The pentoxide is a solid which may be 
sublimed at a high temperature but which gives off almost no 
vapor at ordinary temperatures. It has a very strong affinity 
for water and is the most perfect drying agent for gases which we 
have. If 10,000 liters of air are passed through a comparatively 
small tube filled with the pentoxide, no moisture which can be 
determined remains in the gas, while the vapor of the pentoxide 
which is carried away by the gas weighs only one milligram (Mor- 
ley, Am. J. Sci. 34, 199 (1887) ; J. Am. Chem. Soe. 26, 1171 (1904). 

Another oxide of phosphorus, P2Ox, called phosphorus tetrox- 
ide, is known, but is of little interest except as corresponding to 
nitrogen tetroxide, N2Q4. 

Acids of Phosphorus. While the acids of nitrogen, nitrous 
acid, HNOz, and nitric acid, HNOs, are formed by the addition 
of one molecule of water to the anhydrides, N2O3 and N.Os, 
the normal acids of phosphorus corresponding to these are 
formed by the addition of three molecules of water to the cor- 
responding anhydrides (using the simpler formulas). Two other 
acids, metaphosphoric acid, HPO;:, and pyrophosphoric acid, 
H4P.O;, are also derived from phosphorus pentoxide (phosphoric 
anhydride), P2O;. As the addition of water is considered as 


ACIDS OF PHOSPHORUS 2A] 


neither an oxidation nor a reduction, the three acids derived 
from the pentoxide are all called “‘ phosphoric ”’ acids and are 
distinguished by prefixes. These relations will be clearer from 
the following table: 


(P.0!.3H. O) = 2 H3PO. Hypophosphorous acid 
(P203.38 H,O)= 2 H3PO3; Phosphorous acid 
(P2O;.8 H2O) = 2 H3PO, Orthophosphoric acid 
(P,0;.2 H,O) = HuP,0, Pyrophosphoric acid 
(P.O;.H2O) = 2 HPO; Metaphosphoric acid 


Basicity of the Acids of Phosphorus. The formulas of hypo- 
phosphorous acid, H3PO., phosphorous acid, H3PO3, and ortho- 
phosphorie acid, H3PO,, might lead us to expect each of these 
acids to be tribasic. It is found, however, that only one atom of 
hydrogen in hypophosphorous acid can be replaced by metals 
and only two of the hydrogen atoms in phosphorous acid. In 
other words hypophosphorous acid is monobasic, phosphorous 
acid dibasic, and orthophosphoric acid tribasic. The normal 
sodium salts are: 

Sodium hypophosphite, NaH2POs: 
Sodium phosphite, NazsHPO; 
Sodium orthophosphate, NasPQO, 


These and other facts, which cannot be given here, make it 
probable that the structure of these acids is correctly represented 


by the following formulas : 
HL vy 
Hypophosphorous acid plan : 
H O—H 
H—O O 
Phosphorous acid DRG : 
H O= H 
H—O O 
Orthophosphorie acid Di 
-—O Ot 


1 This oxide is given in many of the books, but its existence is 
extremely doubtful. 


248 A TEXTBOOK OF CHEMISTRY 


According to these formulas only the hydrogen atoms which 
are united to oxygen are acid in character. Also the oxidation 
of the lower acids consists in the introduction of an oxygen atom 
between a hydrogen atom and the phosphorus. According 
to the electron theory the valence of the phosphorus is negative 
toward the hydrogen atoms and positive towards the oxygen. 
Oxidation, in such a case, consists in the change of a negative 
valence to a positive one. 

Hypophosphorous Acid, H3;PO,..° The sodium salt of hypo- 
phosphorous acid, NaH2POs, is formed when phosphorus is 
warmed with a solution of sodium hydroxide, phosphine, PHs3, 
being evolved at the same time. ‘The acid is monobasic. It is 
a powerful reducing agent. Some hypophosphites are used 
in medicine. 

Phosphorous Acid, H;POs;, is formed with phosphoric and 
hypophosphorie acids, when ordinary phosphorus is allowed to 
oxidize slowly in moist air, but it is extremely difficult to separate 
the mixture into its components. The pure acid may be pre- 
pared by the hydrolysis of phosphorus trichloride. It is a bi- 
basic acid, the two sodium salts being monosodium phosphite, 
NaH2PO;, and disodium phosphite, NazgHPO;. Phosphorous 
acid is also a powerful reducing agent. . 

Orthophosphoric Acid, H3POx,, is formed when phosphoric 
anhydride, PsO.o, is dissolved'in hot water. It is also formed 
when solutions of pyrophosphoric acid, H4P20;, or metaphos- 
phorie acid, HPOs, are boiled, especially if some strong acid, as 
nitric acid or hydrochloric acid, is present to catalyze the reac- 
tion, which is to be considered as a hydrolysis : 


OH 

Vinee Oe ange 

re OH 

as + HOH = OH 
O—P—OH a 


Pyrophosphoric Acid Orthophosphorie Acid (2 mols) 


PHOSPHORIC ACIDS 249 


O OH 
o=—pZou + HOH = o—PLon 
Nou 


Metaphosphoric Acid 


Pure orthophosphorie acid forms clear, rhombic crystals, 
which melt at about 40°. These crystals dissolve in a small 
amount of water, forming a heavy, sirupy liquid somewhat 
resembling concentrated sulfuric acid in appearance. 

An impure solution of phosphoric acid was formerly prepared 
on a large scale, as a step in the manufacture of phosphorus, by 
treating bone ash with dilute sulfuric acid and filtering the solu- 
tion from the calcium sulfate, which is only slightly soluble in 
water : 

Ca3(PO,)o +. 3 H.SO, ey H;PO, ote 5 CaSO, 


Tricalcium Calcium 
Phosphate Sulfate 


Orthophosphoric acid forms three classes of salts, in which 
one, two or three atoms of hydrogen are replaced in each mole- 
cule of the acid. These are called primary, secondary and ter- 
tiary, or, more often, mono-, di- and tri-metallic salts. The 
following are the names of the sodium and calcium salts: 


Monosodium phosphate, NaH.PO, (primary) 
Monocalcium phosphate, Ca(H2POx,)2 (primary) 
Disodium phosphate, ' NasHPOs, (secondary) 
Dicalcium phosphate, © CaHPO, (secondary) 
Trisodium phosphate, | Na3PQO, (tertiary) 
Tricalcium phosphate, Ca3(POu.)e (tertiary) 


Orthophosphoric acid is much the most important of the acids 
of phosphorus, being the acid into which all of the others tend 
to pass either by oxidation or hydrolysis. Apart from its oc- 
currence in organic compounds, phosphorus is found almost 
exclusively in the form of orthophosphates, and these phosphates 
are an indispensable constituent of arable soils. 


250 A TEXTBOOK OF CHEMISTRY 


Ionization of Orthophosphoric Acid. Orthophosphoric acid 
is an acid of only moderate strength. A ;5 formular (7.e. con- 
taining one gram molecule in 10 liters of water) solution contains, 
of ‘course, three times as many replaceable hydrogen atoms as 
a tenth normal solution of hydrochloric acid, but it contains only 
one third as many hydrogen ions. This means that in the ioniza- 
tion reaction: 


H;PO, => H+ + H.PO.s 


the equilibrium is comparatively far to the left, even in quite 
dilute solutions. Even in very dilute solutions the second and 
third hydrogen atoms ionize to only a very slight extent. This 
may be either because the three hydrogen atoms in orthophos- 
phoric acid are different or because after the removal of the oe 
hydrogen atom the negative ion, H,POs, holds the remaining 
hydrogen atoms too strongly for them to separate easily as ions. 
The second explanation seems more probable. If the first hy- 
drogen atom is completely neutralized by the addition of a base: 


H3 -b H,PO, +t Nat ole () Ee ee Nat + H,.PO,7 - H.O 


the dihydrogen phosphate ion, H,.PO¢g, will ionize to a slight 
extent: 


HPO; == HPO + Ht, 


but the solution is only faintly acid, and if more sodium hydroxide 
is added to such a solution, the accumulation of the monohy- 
drogen phosphate ions, HPO; shifts the equilibrium to the left. 
This is because sodium salts of weak acids are always much 
more completely ionized than the corresponding acids. ‘This 
shifting of the equilibrium prevents much formation of new — 
hydrogen ions, as those which are present are removed by com- 
bination with the hydroxide ions of the sodium hydroxide. 
Before all of the second hydrogen atoms of the phosphoric acid 
have been neutralized, the tendency of the monohydrogen phos- 
phate ions, HPO4¢, to combine with hydrogen ions will become so 
strong that even the hydrogen ions of water will combine with 
them, leaving an excess of hydroxide ions in the solution. Such 
a solution must, of course, react alkaline. From this conduct 


IONIZATION OF PHOSPHORIC ACID 251 


of phosphoric acid it is evident that if we attempt to titrate a 
solution of phosphoric acid by adding sodium hydroxide, instead 
of the sharp change which occurs in titrating hydrochloric or 
sulfuric acid, there will be a gradual change from a solution con- 
taining a slight excess of hydrogen ions, H*, to one containing 
a slight excess of hydroxide ions, OH~. Two things result from 
these properties of solutions containing salts of phosphoric acid : 
first, unless a very sensitive indicator is chosen, that is, one in 
which the change in color is produced by a very slight change 
in the ratio between the hydrogen and hydroxide ions present, 
the end point of the titration will be indefinite; and, second, 
since most indicators change color, not when the number of 
hydrogen, H*, and hydroxide, OH, are equal, but when 
there is an.excess of one or the other, and this excess 
differs for different indicators, the end point in titrating 
phosphoric acid will depend on the indicator chosen. (See 
p. 387.) Thus, if methyl orange or cochineal is used, the 
end point in fairly dilute solutions will be found when 
the solution corresponds very nearly to the composition 
NaH,PO,. With phenolphthalein, on the other hand, the end 
will correspond nearly to the composition NazHPO.. With 
litmus the end lies between the two. If alizarine green is used, 
the change in color occurs when the composition of the solution 
is very nearly represented by the formula NasPQu. 

It is well, also to consider the conduct of disodium phosphate 
from a somewhat different point of view, which, however, follows 
logically from what has been said. If the salt, which crystallizes 
with the composition NazHPO,.12 H,0, is dissolved in water, we 
should expect the formation of the ions, Nat + Nat + HPO... 
But, as has been stated, in the presence of many of the mono- 
hydrogen phosphate ions, HPOs-, these have a tendency to form, 
with the hydrogen ions of the water, dihydrogen phosphate 
ions, H,PO.-, because the latter ionize only to a slight extent. 
This results in the presence of an excess of hydroxide ions in the 
solution, which will react alkaline toward indicators that are 
sensitive to a slight excess of hydroxide ions. 


202 A TEXTBOOK OF CHEMISTRY 


Nat + Na* + HPO; + H* + OH-— 
Water in 
Ionic Form 


=> Nat + Nat + H,PO, + OH- 


This sort of hydrolysis occurs with all salts of strong bases, 
as sodium hydroxide or potassium hydroxide, with weak acids 
or with acids whose second or third hydrogen atoms undergo 
slight ionization. 

Trisodium phosphate, Na3PO,, will, of course, be much more 
completely hydrolyzed : 


Nat + Nat + Nat + PO= + Ht + OH- 
=> 3 Nat + HPOs + OH- 


The only tertiary salts of orthophosphoric acid which are sol- 
uble in water are those of the alkali metals, sodium, potassium, 
ete. All other tertiary or normal phosphates are insoluble. 
Many of the primary and secondary phosphates are either 
insoluble or are decomposed by water into phosphates which 
approach the tertiary phosphates in composition, and either 
phosphoric acid or more acid phosphates, which dissolve in an 
excess of the acid. 

Decomposition of Primary and Secondary Salts of Ortho- 
phosphoric Acid. Salts of orthophosphoric acid which contain 
hydrogen decompose on heating, losing all of their hydrogen as 
water and leaving salts of metaphosphoric or pyrophosphoric 
acid. As ammonium salts dissociate on heating, these give the 
same products as if they contained hydrogen in place of ammo- 
nium, NH Monosodium phosphate, NaH2PO,, gives sodium 
metaphosphate, NaPO3; and sodium ammonium phosphate, 
NaNHsHPO,, gives the same compound. Disodium phosphate, 
NazHPO,, gives sodium pyrophosphate, NasP2.O;; and am- 
monium magnesium phosphate, NHiMgPO,, gives magnesium 
pyrophosphate, MgoP,0,. Magnesium diammonium phosphate, 
Mg(NH,)4(POx)o, or Mg [(NH4)2POu4]2, gives magnesium meta- 
phosphate, Mg(POs)s. 


PHOSPHORIC ACIDS 253 


Pyrophosphoric Acid, H4P.0;. If orthophosphoric acid is 
heated carefully at 250° it loses water and is changed to pyro- 
phosphoric acid : 


2 H;PO,— HO — H,P.0, 


The acid may be dissolved in cold water, giving a solution which 
differs in its properties from those of a solution of orthophos- 
phoric acid. Especially, after neutralization it gives with silver 
nitrate, AgNOs, a white precipitate of silver pyrophosphate, 
Ag:P207, while orthophosphoric acid, or orthophosphates, 
will give a yellow precipitate of trisilver phosphate, Ag3PQOx. 
Sodium pyrophosphate is easily prepared by heating disodium 
phosphate, Na2HPO.. 

Metaphosphoric Acid, HPOs;, is formed when phosphoric 
anhydride, P40, is dissolved in cold water or when either ortho- 
phosphoric acid or pyrophosphoric acid is heated to a high tem- 
perature. It differs from the other two phosphoric acids in that 
its neutralized solution gives with silver nitrate a white pre- 
cipitate of silver metaphosphate, AgPOs, instead of the yellow 
precipitate of trisilver phosphate, Ag3;PO.,, given by ortho- 
phosphoric acid and the white precipitate of silver pyrophos- 
phate, AgsP.O;, given by pyrophosphoric acid. Metaphos- 
phorie acid also precipitates a solution of albumin, as of the 
white of an egg, which has been acidified with acetic acid, while 
ortho- and pyrophosphoric acids or their salts do not do this. 

Sodium metaphosphate may be prepared by heating either 
monosodium phosphate, NaH»2PQ,, or sodium ammonium phos- 
phate, NaNH,HPO,. This last salt is called microcosmic salt 
and is used in blowpipe analysis: When this salt is heated in 
a loop of platinum wire, it melts to a clear bead of sodium meta- 
phosphate, NaPO3, which will dissolve the oxides of many of 
the metals, forming double salts of orthophosphoric acid : 


NaPO; + CuO = NaCuPOQO, 


Sodium Copper 
Orthophosphate 


254 A TEXTBOOK OF CHEMISTRY 


The metaphosphate may be considered here as, in a certain 
sense, an acid anhydride which with oxides forms normal salts of - 
orthosphoric acid. The copper sodium phosphate is blue, and 
the colors given to the microcosmic bead by different metallic 
oxides serve as a means for their identification. 

* Metaphosphoric acid may be vaporized at a high tempera- 
ture and the vapor has the formula (HPO3)2. <A study of the 
salts and of the properties of solutions of the acid prepared in 
different ways has shown that several polymeric forms of the 
acid exist, that is, forms having the same composition but dif- 
ferent molecular weights. The salts of these various forms are 
called dimetaphosphates, M2P.O,, trimetaphosphates, M3P3Qs, 
tetrametaphosphates, MiP.O12, ete. In these formulas “ M ” 
is used to represent any univalent metal. 

* Hypophosphoric Acid, H.POs, is one of the products formed 
by the slow oxidation of phosphorus in moist air. From the 
solution obtained in this manner the rather difficultly soluble 
acid sodium salt, NaHPOs, is precipitated by a concentrated 
solution of sodium acetate, NaCsH3QOx.. 

From its formula we should expect that phosphorus tetroxide, 
P.O4, would be the anhydride of hypophosphoric acid, but, 
curiously enough, when phosphorus textroxide is dissolved in 
water, a mixture of phosphorous and orthophosphoric acids is 
formed : 


P.O, -|- 3 H.O a H3PO;3 + H3PO, 


The true anhydride of hypophosphoriec acid (PO2?) has not 
been prepared. 

* Sulfides of Phosphorus. Four sulfides of phosphorus have 
been prepared, tetraphosphorus trisulfide, P1453, tetraphosphorus 
heptasulfide, P4S7, triphosphorus hexasulfide, P3S., and diphos- 
phorus pentasulfide, P2S;._ The last is usually called phosphorus 
pentasulfide. It melts at 274°-276° and boils at 530°. It has 
been used in chemical laboratories frequently to obtain a nearly 
constant, high temperature. 

Tetraphosphorus trisulfide, P4S3, melts at 165°-166° and boils 


PHOSPHORUS 250 


at 225°-235° under a pressure of 10 mm. As it takes fire with 
slight friction and as its vapors are either nonpoisonous, or, in 
any case, far less poisonous than those of ordinary phosphorus, 
it is likely to replace the latter entirely for the manufacture of 
matches. 

EXERCISES 


1. Write the equation for the reaction between phosphorus and a 
solution of sodium hydroxide, giving hydrogen and sodium hypophos- 
phite. 

2. Write the equation for the reaction between phosphorus and 
sodium hydroxide in solution, giving phosphine and sodium hypophos- 
phite. 

3. Write the equation for the reaction giving liquid hydrogen phos- 
phide, P2H4, and sodium hypophosphite. 

4, What is the distinction between a substance which sublimes and 
one which boils? Under what conditions does water sublime ? 

5. What percent of phosphorus pentachloride is dissociated when its 
gram molecular volume weighs 156.4 grams? What per cent when it 
weighs 130 grams ? 

6. If on heating phosphorous and hypophosphorous acids the prod- 
ucts formed are phosphine, metaphosphoric acid and water, what are 
the equations representing the decomposition of these acids? Are 
these decompositions consistent with the structural formulas which 
have been proposed for these acids ? 

7. If the structure of phosphorous acid were correctly represented by 


O—H 
the formula PCO—H, how ought it to decompose on heating ? 
O—H 


8. Metaphosphoric acid volatilizes at a very much higher tempera- 
ture than sulfuric acid. What will be.the effect of heating a mixture 
of sodium sulfate and metaphosphoric acid ? 

9. How much iodine and water will be required to convert 10 grams 
of phosphorus into orthophosphoric acid if the reaction is quantitative ? 

10. How many liters of air will be required to burn 10 grams of 
phosphorus to the pentoxide ? 


CHAPTER XV 


ARSENIC, ANTIMONY AND BISMUTH 


Ir has been pointed out that with increasing atomic weights 
the elements of the nonmetallic groups of the Periodic System 
become more metallic in character. This is especially evident 
in the fifth group. Arsenic is metallic in its appearance, opaque 
and like steel on its surface when not tarnished. It is, however, 
brittle, and its chloride, AsCl, is hydrolyzed by water, resembling 
the chlorides of the nonmetals rather than those of the metals. 
Arsenic forms no salts with sulfuric, nitric or other acids. Anti- 
mony and bismuth are still more metallic in their appearance 
and bismuth is malleable to a slight extent. Their chlorides are 
hydrolyzed by water, at first, only to the oxychlorides, SbOCI 
and BiOCl. Both of them form normal nitrates, Sb(NO3)3, and 
Bi(NOs)3, and sulfates, Sbe(SOx.)3, Bix(SO.)3, though these are 
hydrolyzed to basic salts or even to the hydroxides or oxides by 
water. 

Arsenic, As, 74.96. Occurrence. Arsenic is found in the free 
state in nature, but occurs chiefly combined with sulfur, either 
alone, as in the disulfide, realgar, AsoSo, or the trisulfide, orpiment, 
AsoS3, or, much more frequently, with the sulfides of other 
metals, the most common compound of this kind being arseno- 
pyrite, or mispickel, FeAsS. Iron pyrites and copper pyrites 
almost invariably contain arsenic, often in considerable quan- 
tities. From the former the arsenic finds its way into commer- 
cial sulfuric acid and from that into a great variety of chemical 
products. From the copper pyrites the arsenic escapes along 
with the sulfur dioxide in the process of roasting, no less than 
twenty-five tons a day of arsenic trioxide escaping from a single 
smelting furnace in Montana (J. Am. Chem. Soc. 29, 993 (1907). 

256 


ARSENIC Biv 


Preparation and Properties of Arsenic. Metallic arsenic is 
usually prepared by heating arsenopyrite, FeAsS, the arsenic. 
subliming and leaving ferrous sulfide, FeS, behind. Prepared in 
this manner it forms a dark gray, brittle mass. Fragments 
heated in a closed tube or before the blowpipe on charcoal, so 
that the tarnished surface is removed, appear like steel. When 
deposited on a glass or porcelain surface (Marsh’s test), arsenic 
is brown or black according to the thickness of the deposit, 
usually showing brown at the edges where the deposit is thin, 
while antimony is a more sooty black. Here, again, we have an 
increase in metallic properties with increasing atomic weight, 
opacity being one of the most marked properties of metals. 

The specific gravity of gray arsenic is 5.73. The formula of 
its vapor at 560°-670° is Asa, at 1700°, Aso. The melting point 
of arsenic is higher than its boiling point, hence it sublimes with- 
out melting when heated on charcoal or in a tube closed at one 
end, a property which distinguishes it easily from antimony. 

Metallic arsenic is sometimes used for poisonous fly papers. 
Three one-hundredths of a per cent of arsenic lowers the con- 
ductivity of copper 14 per cent and injures it seriously, especially 
for electrical use. 

Arsine, AsH;. Marsh’s Test. When almost any soluble 
compound of arsenic is added to a flask in which hydrogen 1s 
being generated from 
zine and sulfuric or hy- 
drochloric acid, the 
arsenic is reduced to 
arsine. If the hydrogen 
containing arsine is con- 
veyed through a hard 
glass tube, narrowed at 
one point (Fig. 79), and 
the tube is heated just back of the constriction with a Bunsen 
flame, the arsine is decomposed and metallic arsenic is deposited 
as a brown or black mirror on the glass. As small a quantity 
of arsenic as zq/gp of a milligram can be seen in this form, 





Fig. 79 


258 A TEXTBOOK OF CHEMISTRY 


and the process has been long used, under the name of Marsh’s 
test, for the detection and estimation of small quantities of 
arsenic, especially in cases of criminal poisoning, or for the 
examination of wall papers or articles of food. One of the 
first requisites in making the test is, of course, that the zinc, 
sulfuric acid and other materials used should be entirely free 
from arsenic. Commercial zinc and commercial sulfuric acid 
almost invariably contain the element. It is necessary, also, 
to distinguish the mirror from that of antimony, which may be 
obtained in the same manner. When the amount of arsenic 
is considerable, it imparts to the burning hydrogen flame a pale 
blue color, and arsenic is deposited on a piece of porcelain held 
in the flame, very much as soot is deposited from a candle 
flame. 

Arsine may be condensed to a liquid, which boils at — 55°. 
It is very poisonous. Some years ago a chemist in Chile was 
fatally poisoned while working with it. 

Arsine does not show any tendency to combine with acids, as 
ammonia and phosphine do. 

Arsenic “ Trioxide,” AssOg. When arsenic is heated in the 
air, it burns to arsenic “ trioxide,” frequently called white 
arsenic. ‘The simpler formula, As2O3, is commonly used for the 
compound, but the density of its vapor corresponds to the 
formula As,O,. It crystallizes in octahedra which are highly 
characteristic, and the microscopic identification of the crystals 
is one of the most important means of demonstrating the pres- 
ence of arsenic. 

Arsenic trioxide is one of the most common compounds of the 
element and is frequently used as a ratsbane and has often 
been used for criminal poisoning. The fatal dose for an adult is 
from 0.06 to 0.18 gram (one to three grains), but it seems possible 
to accustom the organism to its use, and the so-called arsenic 
eaters may sometimes take four times that amount without 
apparent injury. The best antidote is freshly precipitated 
ferric hydroxide, Fe(OH);, or a colloidal solution of ferric 
hydroxide. 


ARSENIC 259 


Crystallized arsenic trioxide dissolves in 50 parts of water at 
25°. The amorphous form is somewhat more soluble. The 
solution reacts faintly acid, and forms salts with bases, but on 
evaporation it deposits the trioxide. 

Arsenious Acid. As has just been stated, arsenious acid 
resembles sulfurous and nitrous acids in that it exists only in 
solution and decomposes easily into its anhydride and water. 
Salts of acids derived from this anhydride are known, however. 
Among these may be mentioned silver orthoarsenite, Ag3;AsOs;, 
and monopotassium diarsenite, KHAs.O4.. The last seems to 
be derived from a diarsenious acid, H2As,O4, which would cor- 
respond to a doubled nitrous acid, (HNO2)2. Paris green is a 
double salt of copper with acetic and arsenious acids, 
Cu(CsH,0;)2.3 Cu(AsQOz)s. 

Arsenic Pentoxide, As.O;, and Arsenic Acid, H;AsO,. When 
arsenic trioxide is warmed with nitric acid, it is oxidized to 
arsenic acid : 


2 HNO; + As2Oxz -- 2 H,O ray H3AsO, + NO -b NO, 


From a concentrated solution the acid crystallizes with one 
molecule of water, H3AsO4.H.O. At 140°-180° this hydrate 
loses water and gives pyroarsenic acid, H4AseO7, and at 200° 
the latter loses more water and gives metarsenic acid, HAsQOs. 
At a higher temperature metarsenic acid loses more water and 
arsenic pentoxide, As.O;, remains. This cannot be volatilized 
without decomposition, and its true molecular weight is not 
known. Salts of pyroarsenic and metarsenic acids may also be 
prepared by heating secondary and primary salts of arsenic 
acid, but the acids are not known in solution, as they are hydro- 
lyzed by water at once to orthoarsenic acid. ‘Trisilver arsenate, 
Ag3AsOg, is reddish brown and insoluble ; and the white, crystal- 
line magnesium ammonium arsenate, MgNH,AsQy, is also 
insoluble, closely resembling the corresponding phosphate. 

Arsenic acid is an oxidizing agent in concentrated solution, 
liberating chlorine from hydrochloric acid, while it is itself 
reduced to arsenious acid or oxide. But the action is reversible, 


260 A TEXTBOOK OF CHEMISTRY 


and the reverse effect will occur in dilute solutions, chlorine 
oxidizing arsenious acid to arsenic acid: 


H3AsO, + 2 HCl Pras H;AsO3 + Cl, 


In a neutral or faintly alkaline solution the equilibrium is so 
far toward the formation of the arsenate that the oxidation by 
iodine, even, is practically quantitative and is used for the 
standardization of iodine solutions : 


Na3AsQO3 -+- I, 4 Oy NaHCO; = Na3AsO4-+ 2 Nal + H.O -+- 2 CO. 


Arsenic Trichloride, AsCl;, may be prepared by the direct 
union of arsenic and chlorine. It is a colorless liquid, which boils 
at 130°. It is almost completely hydrolyzed by water to 
hydrochloric acid and arsenious oxide or acid. Some arsenic 
trichloride is still present in the solution, however, as a part 
of the arsenic passes over on distilling the solution, while arsenic 
does not escape on distilling a solution of arsenious oxide. 

Sulfides of Arsenic. There are four sulfides of arsenic, 
AsoSe, AseS3, AseSs and As4S3. The last was prepared rather 
recently. 

Arsenic Disulfide, or Realgar, AsoSe, is found in nature and 
may be prepared by melting a mixture of arsenic and sulfur. 
It forms a red, crystalline mass which becomes lighter colored 
when powdered, and was formerly used by artists in painting. 

Arsenic Trisulfide, or Orpiment, As2S3, is also found in nature, 
and is prepared artificially by melting sulfur and arsenic mixed 
in the proper proportion. When prepared in this way it forms 
a yellow, crystalline mass and the powder is used as a pigment, 
especially by artists. From acid solutions of arsenic trioxide 
the trisulfide is precipitated in an amorphous form. It is one 
of the most insoluble sulfides known and is scarcely attacked 
by the most concentrated hydrochloric acid. It is, however, 
dissolved in the presence of oxidizing agents, as by nitric acid, 
aqua regia or potassium chlorate and hydrochloric acid. 

Arsenic Pentasulfide, As2S;, is precipitated from a solution of 
arsenic acid, H3AsO,4, containing hydrochloric acid, apparently 
through the substitution of sulfur for oxygen, giving the series 


ARSENIC 261 


of acids, H3AsSO3, H3AsSeO2, H3AsS30 and H3AsSys. The 
last then dissociates into hydrogen sulfide and arsenic penta- 
sulfide. (MacCay, J. Am. Chem. Soc. 24, 661 (1902); Z. anorg. 
Chem. 29, 36 (1901). It may also be prepared by melting a 
mixture of the elements. 

Sulfarsenites and Sulfarsenates. Arsenic trisulfide, AsoSs, 
and arsenic pentasulfide, AsoS;, dissolve easily in solutions of 
ammonium sulfide, (NH,)2S, or sodium sulfide, NaS, giving 
solutions of sulfarsenites and sulfarsenates . 

5 (NH4)2S “4. ASoS3 =a (NH4)3AsS3 


Ammonium 
Sulfarsenite 


3 NaS ae AsoSs as) NasAsSq 


Sodium 
Sulfarsenate 


These compounds may be considered as arsenites and arsenates 
in which the oxygen has been replaced by sulfur. Antimony 
forms similar compounds, but bismuth does not form them in this 
manner — another illustration of the fact that bismuth is more 
distinctly metallic and does not show the same tendency as 
arsenic and antimony to form acid radicals. The formation of 
these compounds is used in analytical chemistry to separate 
arsenic and antimony from metals which are more metallic in 
character and which do not form similar compounds. 

From solutions of the sulfarsenites or sulfarsenates, acids 
precipitate the arsenic as the trisulfide, AsoS3, or the pentasul- 
fide, AsoSs. 

Colloidal Arsenic Trisulfide. It has been pointed out that 
arsenic trisulfide is one of the most insoluble compounds known. 
It requires at least two million parts of water to dissolve one part 
of the sulfide. In spite of this, however, hydrogen sulfide gives 
no precipitate with a solution of arsenic trioxide in pure water. 
A study of the properties of the solution obtained in this manner 
indicates that the interaction between the trioxide and hydrogen 
sulfide is practically complete : 


As2O3 a 3 HeS:= = AsoSz aa 3 H,O 


262 A TEXTBOOK OF CHEMISTRY 


The solution has the properties of a typical colloidal ‘“solu- 
tion”? — a “solution” in which a substance, which under other 
conditions is insoluble and separates as a precipitate, remains in 
suspension. Such solutions will pass through ordinary filters un- 
changed and under an ordinary microscope they appear to be 
homogeneous. The freezing points and boiling points of such 
solutions are practically identical with the freezing point and boil- 
ing point of the pure solvent —in this case water. ‘This indicates 
that colloids are not in the ordinary molecular condition. The 
ultramicroscope reveals in many colloidal solutions the presence 
of aggregates which have a diameter of from 6 to 60 wu.t Under 
the influence of a considerable electrical potential, colloidal 
arsenic trisulfide moves slowly toward the anode, indicating 
that the aggregates carry negative charges. In the case of 
some colloidal solutions the movement is toward the cathode. 
We may distinguish, therefore, negative colloids, as arsenic tri- 
sulfide, and positive colloids, as colloidal silver. ‘The addition of 
an electrolyte to a colloidal solution will usually cause its precipi- 
tation. In general an electrolyte with bivalent ions, as barium 
chloride, BaCls, is more effective than one with univalent ions, 
as sodium chloride, but the effect is dependent also on the degree 
of ionization of the electrolyte and it seems to be the cation 
(e.g. Bat*) which is effective in precipitating a negative colloid, 
and the anion (e.g. Cl” or SOs-~) which precipitates a positive 
colloid. The cation (or anion) is retained by the precipitate 
and cannot be washed away, though it may be displaced by 
another ion of the same sign. 

These facts are, at present, best understood in the light of 
the following theory. In the colloidal solution aggregates of a 
substance which is usually insoluble are formed around negative 
or positive ions, forming in the first case negative, in the second 
case positive, colloids. These aggregates are very much larger 
than ordinary molecules, but they are prevented from falling to 


+ stands for one micron, qogy of a millimeter. uu stands for 
roo Of a micron or one millionth of a millimeter. The wave 
length of sodium light is approximately 0.6 u. 


ANTIMONY 263 


the bottom of the solution, partly because they are still very 
small, but more because on account of their electrical charges 
they are prevented from cohering with other similar aggregates 
to form larger particles and also because there must always 
be in the solution, to balance them electrically, other, ordinary, 
ions with charges of the opposite signs. If these aggregates 
were to separate from the solution as a precipitate, the solution 
would be electrically positive and the precipitate negative in the 
case of arsenious sulfide. When an electrolyte, as barium 
chloride, is added to such a solution, the positive barium ions, 
Ba*tt, combine with the negative ions of the colloidal arsenious 
sulfide, forming neutral aggregates which can then cohere to 
larger aggregates and form an ‘ordinary precipitate. At the 
same time the negative chloride ions, Cl~, balance the positive 
ions of the solution, usually hydrogen ions, H*, and the solution 
remains electrically neutral although the colloid has separated 


from it. It has been shown that in such a case the solution 


remains acid in exact proportion to the amount of barium carried 
down by the precipitate, and it has already been pointed out 
above that the barium cannot be removed from the latter by 
washing. 

A knowledge of the conditions which govern the formation and 
precipitation of colloids is often of very great importance in 
analytical chemistry. The phenomena of colloidal solutions 
also play a very important part in the digestion and assimilation 
of food and in the life processes of both plants and animals. 

Antimony, Sb, 121.77. Occurrence and Preparation. Small 
quantities of antimony are found free in nature, but the element 
occurs chiefly in the mineral stibnite, antimony trisulfide, Sb2Ss3. 
When this is heated in the air, the sulfur burns away as sulfur 
dioxide, SO2, and the antimony remains as the tetroxide, Sb2Ox. 
The process is called roasting and is a very common method of 
treating ores which contain sulfides of the metals. The crude 
oxide is then reduced by heating it with coke or charcoal and 
suitable substances to form a fusible slag with the impurities of 


the ore: Sb.O, + 2C = 2Sb+ 2 CO, 


264 A TEXTBOOK OF CHEMISTRY 


Properties. Antimony is a silver-white, brittle, crystalline 
metal. The specific gravity is 6.52... It melts at 630° and boils 
at 1300°. 

A curious form of the element known as explosive antimony 
can be prepared by the electrolysis of a solution of antimony 
chloride, SbCls, in hydrochloric acid. It has a specific gravity of 
only 5.78. When rubbed in a mortar or when the dry sub- 
stance is heated to 200°, it explodes violently, with an appearance 
of light and heat, being converted into ordinary antimony. The 
transformation is accompanied by the evolution of about 
20 calories per gram. 

When antimony is heated in the air, on charcoal, it melts easily, 
differing in this respect from’ arsenic, which sublimes without 
‘ melting. It burns slowly, giving vapors of antimony trioxide, 
Sb203. It does not dissolve in hydrochloric acid, but is easily 
converted into a mixture of insoluble oxides by nitric acid. 
Antimony and tin are the only metals which are acted upon by 
nitric acid in this manner, giving insoluble oxides or acids. 

Uses. Metallic antimony is a constituent of many important 
alloys, especially of type metal (lead, tin and antimony), stereo- 
type metal (lead, tin, antimony and bismuth), britannia metal 
(tin and antimony) and antifriction metals (lead, antimony and 
copper with a little bismuth) used for bearings in machinery. 
In type metal it gives hardness to the alloy and also causes it to 
expand slightly as it solidifies in the mold, giving clear-cut 
type. 

Stibine, SbHs, is formed when a soluble compound of antimony 
is Introduced into a hydrogen generator containing zine and 
hydrochloric or sulfuric acid. It resembles arsine closely, but 
gives a somewhat more sooty spot on porcelain or in a glass tube 
by Marsh’s test. It is decomposed into antimony and hydrogen 
at a lower temperature than arsine, the decomposition taking 
place slowly at ordinary temperatures, especially in the presence 
of metallic antimony. For the methods of distinguishing 
between the deposits of arsenic and antimony, works on ana- 
lytical chemistry should be consulted. 


ANTIMONY 265 


Oxides of Antimony. Antimony forms three oxides: anti- 
mony trioxide, Sb4O, (or Sb2O3), formed by burning antimony 
in the air or by heating the hydroxide, Sb(OH)3; antimony 
tetroxide, Sb.O,, formed when either the pentoxide or the 
trioxide is heated with free access of air; and antimony pentox- 
ide, Sb2Os, obtained by repeated evaporation of metallic anti- 
mony or one of the lower oxides with nitric acid and finally 
heating the residue to 300°. At a higher temperature it is 
decomposed into the tetroxide and oxygen. ‘The trioxide is the 
only oxide which can be converted into a vapor without decompo- 
sition, and so is the only one for which we really know the molec- 
ular weight and true formula, SbsO,._ It is altogether probable 
that the tetroxide and pentoxide have more complex formulas 
than those given, and, indeed, it is quite possible that the solid 
trioxide has a higher molecular weight and more complex formula 
than that of its vapor. For inorganic compounds, in general, 
it is more convenient to use the simplest formulas which express 
the composition in whole atomic weights. For this reason chem- 
ists continue to use the formulas P2203, P2O;, Sb2O3, ete., and the 
corresponding names, for compounds whose molecules are known 
to be more complex. When questions of structure are con- 
sidered, however, it is important to remember that the molecules 
are more complex than these formulas indicate. 

Antimony hydroxide, Sb(OH)3 or H3SbO3, Antimonious Acid, 
may be prepared by the precipitation of a solution of tartar 
emetic (see below) with dilute sulfuric acid. As is to be expected 
from the position of antimony in the Periodic System, it is 
amphoteric in character, that is, both an acid and a base. Ina 
solution of a strong base it gives up hydrogen and forms a salt 
in which it furnishes the acid radical : 


H;SbO; + NaOH = NaH2SbO;.H:0 


Sodium 
Antimonite 


In a solution of a strong acid, on the other hand, it gives up 
its hydroxyl and forms salts in which the antimony is the metallic 


element: 9 Sh(OH)s + 3 H2SOx = Sbo(SOx)s + 6 H20 


266 A TEXTBOOK OF CHEMISTRY 


In further agreement with this character, the hydroxide or 
acid loses both hydrogen and hydroxyl (OH) easily even when in 
contact with water, going. over into the oxide, Sb2O3: 


As is to be expected, also, both classes of salts are hydrolyzed 
by water. The salts of the alkalies react strongly alkaline in 
solution, while the salts in which antimony forms the metallic 
part are mostly decomposed by water with the precipitation of a 
basic salt and liberation of the free acid. In these basic salts 
| Yeti ini 8 
instead of the group Pe ii which might be expected, the 


pL? 
group Sb , antimonyl, formed from this by the loss of hydro- 


gen and hydroxyl, is often present : 


Antimonyl 
Sulfate 


Tartaric Emetic, KSbOC,H4O,. One of the most interesting 
and important basic salts of antimony is tartar emetic, or potas- 
sium antimonyl tartrate. Cream of tartar, or acid potassium 
tartrate, KHC4H,Og, is the acid potassium salt of tartaric acid, 
H.2C4H4Og, an acid found in the juice of grapes. When antimony 
trioxide, Sb2O3, is boiled with a solution of cream tartar, it dis- 
solves, forming tartar emetic: 


Sb.03 + 2 KHC,H.0, a 2 KSbOC,H105¢ + H.O 
In the tartar emetic the univalent antimony] group, BS F 


replaces hydrogen as though it were a univalent metal, very 
much as the ammonium group, NHg, replaces hydrogen in the 


ANTIMONY 267 


formation of ammonium salts. Tartar emetic dissolves easily 
in water. It is sometimes used as an emetic. 

Antimonic Acids. ‘Three antimonic acids are known, cor- 
responding to the phosphoric acids of similar formulas : metanti- 
monic acid, HSbOs, pyroantimonic acid, H4Sb2O7, and orthoanti- 
monic acid, H3SbO4. Very few salts of the last are known. 

Chlorides of Antimony. Antimony forms three chlorides. 
Antimony trichloride, SbCl;, can be prepared by dissolving the 
trioxide, Sb2O3, or the trisulfide, SbeS3, in concentrated hydro- 
chlorie acid, evaporating the solution and distilling the residue. 
It is a solid which melts at 73.2° and boils at 223°. It is decom- 
posed by water with the precipitation of the oxychloride, which 
may be called antimonyl chloride, SbOCl. From the method 
of preparing the trichloride it is evident that the decomposition 
by water is a reversible reaction, the direction of which depends 
on the concentration of the reacting substances. 

* Antimony tetrachloride, SbCl:, and Hydrotetrachloroanti- 
monic acid, H,SbClg. When solutions of antimony trichloride, 
SbCl3, and antimony pentachloride, SbCl;, in hydrochloric acid 
are mixed, a dark brown solution is formed. The depth of 
color increases on warming and decreases on cooling, indicat- 
ing that the tetrachloride is formed with an absorption of heat, 
since its formation is promoted by an increase of temperature 


(p. 111): SERCH. Chl, 2. SbCl, 


The tetrachloride has not been prepared in pure condition, but 
double salts with other metals such as the cesium tetrachloro- 
antimonate, CseSbCl,, have been prepared. This corresponds 
to an acid, HeSbCls, which probably exists in the dark brown 
solution referred to above and which may be called hydrotetra- 
chloroantimonie acid. 

* Antimony Pentachloride, SbCl;, is formed when antimony 
is burned in chlorine. It is a white solid at low temperatures, 
but melts at — 6° and boils at 140°. At the latter temperature 
it dissociates, partly, into the trichloride and chlorine, exactly 
as phosphorus pentachloride, PCls, does. 


268 A TEXTBOOK OF CHEMISTRY 


* Metachloroantimonic Acid, HSbCl..44 H.O. If chlorine 
is led into a concentrated solution of antimony trichloride in 
hydrochloric acid till the solution becomes colorless or light 
yellow and the solution is evaporated in a current of hydro- 
chloric acid to prevent hydrolysis, very hygroscopic crystals of 
metachloroantimonic acid can be obtained. This may be con- 
sidered as metantimonic acid, HSbOs, in which the three oxygen 
atoms have been replaced by six chlorine atoms. The freezing 
point of the solution indicates that the compound separates into 
the ions H*and SbCl¢. The solution gives a precipitate with 
silver nitrate only after some time, indicating that very few 
chloride ions, Cl”, are present. 

Very many salts of this acid have been prepared, among which 
the following may be mentioned : 


Potassium metachloroantimonate . . . . KSbCl..H.O 
Calcium metachloroantimonate . . . . Ca(SbCle)o.9 H.O 
Aluminium metachloroantimonate . « Al(SbCle)s, loss 


Antimony Trisulfide, Sb2S3, is obtained as an orange-red pre- 
cipitate when hydrogen sulfide is passed into an acid solution of a 
salt of antimony. It also occurs in nature as the black mineral, 
stibnite. It dissolves readily in concentrated hydrochloric acid, 
differing very markedly from arsenic trisulfide in this regard. 

Antimony Pentasulfide, SboS;, is best obtained by the decom- 
position of sodium sulfantimonate with hydrochloric acid : 


2 NasSbS, + 6 HCl =.6 NaCl + Sb.S + 3 HS 


Sulfantimonites, M3;SbS;3, and Sulfantimonates, M3SbS., may 
be obtained in the same manner as the corresponding sulfar- 
senites and sulfarsenates (p. 261). The alkali salts are soluble 
in water, hence the sulfides of antimony dissolve in solutions 
of sodium sulfide or ammonium sulfide and may be separated 
in this way from the sulfides of elements which are very decidedly 
metallic in character. 

Bismuth, Bi, 209. Occurrence, Properties, Uses. Bismuth 
is less abundant in nature than arsenic or antimony, as is usually, 
though not invariably, the case with elements of high atomic 


BISMUTH 269 


weights. It is found mostly in the free state, but is found also 
as the sulfide, Bi2S3, both alone and with other sulfides, espe- 
cially with lead sulfide. It is also found as the oxide, BiQs. 
It is obtained commercially as a by-product in the electrolytic 
refining of lead. The specific gravity of the distilled metal is 
9.78. The melting point is 271°. As bismuth expands on solidi- 
fying, the melting point is lowered by pressure (Principle of 
van’t Hoff-Le Chatelier, p. 111). Its boiling point is below 
1700°, but is not accurately known. 

Bismuth is used in a variety of alloys, usually because it lowers 
their melting points and renders them more suitable for specific 
purposes. It is used in this way in stereotype metal to give an 
alloy which can be cast ina papier-maché mold without injuring 
it, in Wood’s metal (Bi, 4 parts, Pb, 2 parts, Sn, 1 part, Cd, 1 part), 
which melts at 60.5°, and is used for heating baths in chemical 
laboratories, and for many other easily fusible alloys used as 
safety plugs in steam boilers and in automatic sprinklers for 
protection against fire, also for safety fuses in electrical work. 
The addition of a little bismuth has been found an advantage in 
Babbitt metal and in other antifriction metals used for bearings 
in machinery. 

Oxides of Bismuth. Bismuth forms two well-characterized 
oxides, BiO and Bi.O3. It also forms one or more higher oxides, 
called peroxides, for which we should expect the formulas Biz:O, 
or Bi,O;. These higher oxides do not seem to have been obtained 
in a state of purity, probably because of the ease with which 
they and their hydrates lose water and oxygen. A mixture of 
these oxides or hydrates containing some sodium is prepared 
and has been called, without good reason, sodiwm bismuthate. 
It oxidizes manganese compounds to permanganic acid, HMnOQ,, 
in nitric acid solutions and is used for that purpose in the de- 
tection and quantitative determination of manganese. Appar- 
ently no pure bismuthic acid or salt of bismuthic acid has been 
prepared. 

Bismuth Chloride, BiCl;, is formed by the direct union of 
bismuth and chlorine or by the solution of bismuth trioxide, 


270 A TEXTBOOK OF CHEMISTRY 


Bi,O3, in concentrated hydrochloric acid. It melts at 225°-230° 
and boils at 427°-429°. It dissolves in moderately strong 
hydrochloric acid to a clear solution, but the addition of water 
causes the precipitation of bismuth oxychloride, or bismuthyl 
chloride, BiOCI, which is extremely insoluble. 

Bismuth Nitrate, Bi(NO3)3.5 H.O, can be prepared by dis- 
solving either bismuth trioxide, Bi.O03, or metallic bismuth in an 
excess of nitric acid and evaporating to crystallization. When 
bismuth nitrate is treated with water, it is hydrolyzed with the 
formation of a mixture of basic nitrates, which varies in compo- 
sition according to the method by which it is prepared. The 
simplest of these compounds are Bi(OH)2.NO3 and BiONOs: 


OH 
Bi(NO;)3 + 2HOH = Bi<OH + 2 HNO; 
‘NO; 


The mixture of basic nitrates is called in many medical works 
“bismuth. subnitrate,” an antiquated name which does not 
correspond to modern scientific usage. It is used in medicine 
and also as a slightly antiseptic face powder. 

Bismuth Trisulfide, Bio.S3, separates as a black or brownish 
black precipitate when hydrogen sulfide is passed into a solution 
of a soluble bismuth salt. Owing to the more metallic character 
of bismuth, it does not dissolve appreciably in solutions of 
sodium sulfide or ammonium sulfide as the sulfides of arsenic 
and antimony do. It dissolves easily in warm nitric acid, form- 
ing bismuth nitrate, Bi(NQO3)3. 

The following tables of the more important compounds of the 
elements of the fifth group will be of service in reviewing and 
comparing these. Compounds which correspond for different 
elements are selected, especially, for the table. Many other 
compounds are, of course, known. 


ARSENIC, ANTIMONY AND BISMUTH 271 


















































Oxides 
N, 14 Poor As, 75 Sb, 120 Bi, 208 
N2O roe 
NO Saas —— —_—— BiO 
N.O3 P40. AssOg¢ Sb4O¢ Bi,O; 
NOz, N20, —— a SbeO4 
NO; P4O 40 AsO; Sb20s5 Bi.O; ? 
Chlorides 

ee ates fll La BiCl, 
NCls PCl; AsCl3 SbCls BiCls 

SbCl, 
— PCI; a aa SbCl; a amg 

Acids 
H.N2O2 H;PO, —— Se 
HNO, H3PO3 H3AsQO3 H3SbO3 (Bi(OH)3) 
—— H,PO3 
HNO; HPO; HAsO; HSbO; HBi0 3? 
ae H,P.0, H,As,O;, H,Sb.0, 
= oa H3PO, H3AsQO, H3SbO, Yas ae 
Salts of Sulfur Acids 
ee Na3sPS3 Na3AsS3 NagSbS3 NaBis. 
er M4P.2S, Na3AsSq NazSbS4 
Hydrides 

NH; PEs AsH3; SbH; a ian 
NeH, PoH, on 
N3H PioH, =a ae aa es 


Vanadium (V, 51.0), Columbium (or Niobium) (Cb, 93.5), 
Tantalum (Ta, 181.5). These elements, which are found in the 
fifth group in the alternate rows of the Periodic System, are 


272 A TEXTBOOK OF CHEMISTRY 


more decidedly metallic in their properties, corresponding to 
their positions in the system ; and while they show many analogies 
with the elements described in this chapter, they will be reserved 
for a later consideration (p. 522). 


EXERCISES 


1. What are the reactions in Marsh’s test, if arsenious oxide is used ? 
What, if arsenic acid is used ? 


2. If an arsenic mirror on porcelain is warmed with ammonium 
sulfide, it dissolves and the addition of hydrochloric acid gives a lemon- 
yellow precipitate which does not dissolve in concentrated hydrochloric 
acid. Write the equations. 


3. The antimony mirror conducts itself in a similar manner, but 
gives anorange precipitate with dilute hydrochloric acid which dissolves 
in concentrated hydrochloric acid. Write the-equations. 


4, What volume of air will be required for the complete combustion 
of one volume of arsine ? 


5. What weight of arsenious oxide will be required to give one 
pound (453 grams) of Paris green ? 

6. What will be formed by the ignition of magnesium ammonium 
arsenate ? 


7. The specific gravity of fused arsenic trisulfide is 2.76. If a particle 
of the colloidal sulfide has this specific gravity and is one micron in 
diameter, how many such particles would there be in one gram of the 
sulfide? Assuming that it takes 1.5 x 107! molecules of arsenious sul- 
fide (As2S3) to weigh one gram, how many molecules would there be 
in such a particle of colloidal arsenic trisulfide ? 


8. Assuming the formula Bi,O; for bismuth peroxide, what is the 
equation for the reaction between this compound and manganese ni- 
trate, Mg(NOs3)s, in the presence of dilute nitric acid ? 


CHAPTER XVI 


CARBON 


Carbon, C, 12. Occurrence. Although carbon forms only 
about one five-hundredth part of that portion of the earth which 
we can examine, it is in many respects the most important of all 
of the elements. It forms an indispensable element in all living 
organisms, both animal and vegetable, and, indeed, it seems to be 
the peculiar properties of carbon rather than those of any other 
element, which make life, as we know it, possible. In addition 
to this preéminent réle in living bodies, carbon is the principal 
constituent in all substances used for fuel and is the element 
by means of which iron is reduced from its ores. 

The unique character of carbon is suggested by its position 
in the Periodic System. It stands in the first row midway 
between the most strongly nonmetallic element fluorine and one 
of the alkali metals, lithium. Corresponding to this position 
its valence is four, but it combines both with the positive hy- 
drogen and with the negative fluorine and chlorine. Even in 
its elementary forms, it is nonmetallic, transparent and a 
nonconductor of electricity in the diamond, but approaches 
the metals in being opaque and a fairly good conductor of elec- 
tricity in graphite. This double character seems to be closely 
connected with the power which carbon atoms have to combine 
with each other as well as with other elements. The compounds 
of the element are bewildering in their variety and complexity. 
More than 100,000 such compounds have been prepared and 
analyzed, and some thousands of new compounds are discovered 
every year. On account of their number and many peculiarities, 
which distinguish them from the compounds of other elements, 
but also because of their importance and because so much time 


has been devoted to their study, the compounds of carbon are 
273 


274. A TEXTBOOK OF CHEMISTRY 


usually considered separately as a special subject, called organic 
chemistry. 

The study of these compounds has proved so important in 
its relation to the problems of general chemistry, however, that 
no textbook of inorganic chemistry 1s complete without a de- — 
scription of some of them. 

Diamonds. A diamond of the first quality, weighing a carat, 
or 0.2 gram, when properly cut, is worth approximately $500, 
while a kilogram of carbon in the form of coal or coke is worth 
less than one cent. This fact has been a constant challenge to 
chemists ever since the composition of diamonds was discovered. 
It was not, however, till near the close of the nineteenth century 
that even microscopic diamonds were prepared artificially ; and 
even now the theoretical conditions for their preparation are 
not fully understood and no one has succeeded in making dia- 
monds large enough to be of commercial value. 

In 1892 Friedel discovered that a meteor which had fallen in 
Canon Diablo, Texas, contained carbon. By a eareful examina- 
tion, Moissan, in Paris, demonstrated the presence of microscopic 
diamonds in the material. As he stated it. afterwards, nature 
had been caught in the act of making diamonds. The high 
specific gravity of the diamond in comparison with graphite seems 
to have suggested to Moissan that diamonds were probably 
formed under conditions of high pressure. This and the dis- 
covery of diamonds in the meteorite suggested to him the follow- 
ing experiment. A mass of iron was heated to a very high tem- 
perature, 3000°-3500°, in an electric furnace and some pure 
sugar charcoal was dissolved in the hot iron. The mass was 
then suddenly thrust into water. This caused the exterior 
surface to solidify while the interior was still fluid and intensely 
hot. As the interior cooled, some of the dissolved carbon sepa- 
rated in crystalline form, and, as iron containing carbon ex- 
pands on solidifying, the interior portions were subjected to an 
enormous pressure. The lowering of the melting point of the 
iron by the pressure may also have had something to do with the 
success of the experiment. When cold, the iron was dissolved 


DIAMONDS 21D 


in acids, and silicon, graphite and other substances were removed 
by oxidation, solution, and finally by treatment with a liquid 
having a specific gravity greater than that of graphite and less 
than that of the diamond. A few minute crystals were dis- 
covered which were heavier than this liquid. By carefully 
rubbing one of the crystals against the surface of a ruby, it was 
shown that the surface of the latter was scratched. Only the 
diamond and carborundum are known to be harder than the 
ruby, and the specific gravity of carborundum is less than that 
of the artificial diamonds. After securing enough of the crystals 
to weigh a few milligrams, they were burned in a current of 
oxygen and the carbon dioxide formed was absorbed and weighed. 
Twelve parts by weight of the crystals gave 44 parts by weight 
of carbon dioxide. The crystals consisted, therefore, of pure 
carbon and were in reality diamonds. 

Until the beginning of the eighteenth century diamonds had 
been found only in India. In 1727 they were discovered in 
Brazil, at the beginning of the nineteenth century in the Ural 
mountains, and in 1867 at Kimberley in South Africa. The 
Kimberley mines now furnish most of the diamonds for the 
world’s market. The product of the mines is valued at about 
$15,000,000 annually. 

The diamond crystallizes in octahedra and cubes of the iso- 
metric system. Its specific gravity is 3.5. It is the hardest 
known substance, and is not dissolved or oxidized by any known 
liquid or gas at ordinary temperatures. Its index of refraction 
is extraordinarily high, being 2.417 for sodium light. This and 
the high dispersive power give to diamonds, which are cut so 
as to accentuate these properties, the ability to reflect and re- 
fract light in such a manner as to produce brilliant colors. 

For cutting glass, an edge produced by cleavage must be used. 
For diamond drills to be used in boring in rocks in such a manner 
that a solid core can be removed, inferior black diamonds, called 
carbonado, set into the ends of tubes, are employed. 

When heated to a high temperature with exclusion of air, the 
diamond is changed to graphite. When heated in oxygen, carbon 


276 A TEXTBOOK OF CHEMISTRY 


dioxide begins to be formed at 720° and the diamond takes fire 
and burns at 800°-850°. 

Graphite, asecond crystalline form of carbon, isfound in nature, 
especially in Ceylon, Siberia, England and Canada. It may be 
prepared artificially by crystallizing carbon from cast iron, 
from one to two per cent of graphite being left behind when gray 
cast iron is dissolved in acids. It is also formed when any form 
of carbon is heated to a very high temperature in an electric 
furnace. The presence of silicon, aluminium, iron and other 
elements seems to assist in the transformation by the interme- 
diate formation of carbides. In the Acheson process, which has 
acquired considerable technical importance, impure carbon, as 
coke or anthracite, is found more suitable than pure carbon. 

Graphite crystallizes in six-sided leaflets of the monoclinic 
system. It is soft and has a gray, metallic luster and gives a 
metallic streak. Its specific gravity when pure is 2.255. In 
oxygen it begins to give carbon dioxide at 570° and takes fire 
at 690°. At very high temperatures it seems to be the most 
stable form of carbon, into which all other forms tend to pass. 
The total energy of graphite at ordinary temperatures is, how- 
ever, greater than that of the diamond at ordinary temperatures, 
as is evident from the following table of heats of combustion, as 
determined by Berthelot : 


12 grams of diamond give 94,310 small calories. 
12 grams of graphite give 94,810 small calories. 
12 grams of amorphous carbon give 97,650 small calories. 


From this table it appears that if 12 grams of graphite could 
be changed to diamond at ordinary temperatures, 500 small 
calories would be evolved. | 

Graphite is used for “lead”’ pencils, as a lubricant, especially for 
surfaces of wood and where bearings are subjected to a high tem- 
perature, and in making crucibles for use at very high tempera- 
tures for melting and casting steel and difficultly fusible alloys. 
For the last purpose it must be mixed with some fire clay, 
which binds the particles of graphite together and also protects 


CARBON ar 


the graphite from burning by giving an incombustible surface. 
Graphite is also used as a lubricant in a colloidal solution in 
oil (Acheson). It is used for stove polish, to protect iron from 
rusting. 

Amorphous Carbon. When almost any compound of carbon is 
heated, it will decompose with the separation of charcoal or car- 
bon. It is, however, extremely difficult to obtain perfectly pure 
carbon in this manner. The purest carbon is obtained by heat- 
ing sugar, which contains only carbon, hydrogen and oxygen, to a 
high temperature, but it seems doubtful whether the last traces of 
hydrogen can be expelled without the use of a temperature which 
would convert the amorphous carbon partly into graphite. The 
hydrogen may be almost completely removed, however, by 
heating the charcoal in a current of chlorine at 1000°. There 
seems to be no form of amorphous carbon which can be properly 
spoken of as a definite chemical individual, as the density and kin- 
dling temperatures and conductivity for electricity vary gradually 
from a form which has a density of 1.45 and a kindling tempera- 
ture of 300°, up to forms which approach closely to graphite in 
their properties. This fact is doubtless intimately connected 
with the almost infinite variety of ways in which carbon atoms 
unite with each other in the compounds of carbon. 

A great variety of impure forms of amorphous carbon are 
known. All of these contain at least some hydrogen and most 
of them contain oxygen and other elements. 

Lampblackis the soot deposited from substances rich in carbon, 
such as naphthalene, rosin, petroleum, etc., burning with a smoky 
flame. It is used as a pigment, especially in printers’ ink. The 
insoluble and indelible quality of carbon makes such ink even 
more permanent than the paper on which it is printed. 

Wood Charcoal is manufactured by piling wood in heaps, 
covering it with sod and setting fire to it in such a manner that 
only a portion of the wood burns while the heat converts the re- 
mainder into charcoal. The process is wasteful and has been 
largely replaced by methods of charring in retorts or chambers 


‘so arranged that the wood tar, wood alcohol and acetic acid, 


278 A TEXTBOOK OF CHEMISTRY 


which are formed by the decomposition of the wood, may be 
recovered. Combustible gases, which are also formed, are util- 
ized to heat the retorts or chambers. Charcoal is used by tinners, . 
in small charcoal furnaces, for filtering alcohol to purify it and 
for the manufacture of a high grade of iron. It was formerly 
used in very large quantities for this last purpose, but has been 
almost entirely displaced by coke and coal. 

Charcoal, because of the infusible character of all forms of car- 
bon, retains the original structure of the wood and contains an 
immense number of pores of microscopic size. Apparently for 
this reason freshly ignited charcoal will absorb many times its 
volume of gases, especially of those gases which are easily lique- 
fied, such as ammonia, or hydrogen sulfide. This phenomenon 
is called adsorption, and seems to be due to the condensation of 
the gas on the very large surface offered by the porous charcoal. 
Charcoal cooled by liquid air is a very efficient means for ab- 
sorbing residual gases and producing a high vacuum. 

Charcoal was formerly much used in domestic water filters, 
but it has been found that such filters are only very temporarily 
effective. 

Animal Charcoal and Bone Black are obtained by charring 
the refuse of slaughterhouses and bones. They show the prop- 
erty of absorption, especially for coloring matters, in a very 
high degree, and are used in the removal of color from sirups 
and other liquids, as in the purification of sugar. They are also 
frequently used for the purification of organic compounds in 
chemical laboratories. For such use they should be purified by 
treatment with acids to remove calcium phosphate and other 
mineral matters which they contain. 

Coke bears very much the same relation to bituminous coal 
that charcoal does to wood. It is still manufactured in America, 
chiefly in the so-called ‘“ beehive’? ovens — hemispherical 
chambers built of brick, 12 feet in diameter and seven and one- 
half feet high. These are charged with coal while still hot from 
a previous charge, and the volatile matter given off from the 
coal takes fire and burns within the oven, over the surface of the 


CARBON 279 


coal, furnishing the heat necessary to convert the coal into coke. 
An opening on one side of the oven supplies air and the products 
of combustion escape through a circular opening at the top. 
This wasteful method is being slowly replaced by methods of 
coking in retorts, by the use of which it is possible to recover the 
tar, ammonia and combustible gases. The gas produced is 
formed in excess of what is necessary for heating the retorts, 
and may be used in part for other purposes, while in the beehive 
ovens nothing is saved except the coke. 

Coke is used chiefly in blast furnaces for the production of cast 
iron. It is used in some other metallurgical processes and to a 
limited extent as a domestic fuel. 

Gas Carbon. Carbon Electrodes. On the walls of the retorts 
used in the manufacture of illuminating gas, carbon is deposited 
from the decomposition of carbon compounds in the volatile 
matter given off by the coal. As a result of the prolonged 
heating it assumes a semicrystalline, dense form, approaching 
graphite in its properties. It has a density of 1.9-2.0, a very 
high kindling temperature, and is a fairly good conductor of 
electricity. This gas carbon, or frequently, also, anthracite, 
petroleum coke, or some other form of amorphous carbon, is 
mixed with a little coal tar or some petroleum product for a 
binding material and molded into various forms for use as 
electrodes in the electroylsis of sodium chloride, aluminium com- 
pounds or other substances, for use in electrical furnaces and for 
the carbon electrodes of are lights. The mixture is subjected 
to a hydraulic pressure of 500 atmospheres to.render it as dense 
as possible, and is then heated to a temperature of 1200°-1400° 
for 24-48 hours, till all volatile compounds have been expelled 
and the carbon has become dense and hard and a good electrical 
conductor. Such electrodes are scarcely attacked by the chlo- 
rine evolved in the electrolysis of a solution of sodium chloride 
and are scarcely affected when heated, out of contact with air, to 
any temperature below that of the electric arc. Their resistance 
to oxidizing agents in electroylsis may be further increased by con- 
verting them partly or wholly into graphite in an electric furnace. 


280 A TEXTBOOK OF CHEMISTRY 


Coal. Very much the same process which occurs rapidly 
when wood is heated seems to have gone on slowly through some 
hundreds of thousands or millions of years with vast quantities 
of woody material accumulated during certain periods of geolog- 
ical time and afterwards covered with thick layers of clay and 
other materials, beneath the surface of ancient oceans. Wood 
consists chiefly of carbon, hydrogen, oxygen and nitrogen, with 
small quantities of mineral matter. The transformation to 
coal has been occasioned by the gradual loss of oxygen and some 
of the hydrogen, in such a manner that the per cent of carbon 
gradually increases. ‘The per cent of the oxygen decreases until 
that element nearly disappears in anthracite. The per cent of 
hydrogen decreases only slightly till the last stage — the trans- 
formation to anthracite —is reached. It seems probable that 
this last transformation occurred at a more elevated temperature. 
The changes in composition which have taken place during these 
transformations are apparent in the following table: 


CHANGES OF Woop MATERIAL DURING GEOLOGICAL TIME! 























PERCENTAGE COMPOSITION 
EXCLUSIVE OF MOISTURE CaLor ES 
F PERCENT- | PERCENT- VALUE; 
MatTERIAL AND ARE AGE OF AGE OF CALORIES 
Care u ftydenst Oxyoul nies ASH Moistur# | PER KIL- 
bon gen n en. pores 
g ge ge 
Wood — Oak .| 50.35 6.04 43.52 0.09 0.37 20.00 2 3696 
Peat . | 99.70 5.70 33.04 1.56 11.84 14.24 2 3979 
Brown Lignite, 
North Dakota| 74.88 4.99 19.12 1.01 9.35 35.38 3846 
Black Lignite, 
Colorado . 76.83 5.34 16.29 1.54 5.99 18.68 5636 
Bituminous, 
Illinois 83.42 5.29 9.52 ied, 11.28 8.50 6542 
Semibitumin- } 
ous, West Vir- 
ginia Poca- 
hontas . 91.50 4.38 3.07 1.05 6.55 3.67 7939 
Anthracite 93.76 2572 oy i 0.41 10.80 2.18 7216 
Charcoal . 84.11 1706 14.36 _ 2.50 —- 6626 
Coke 95.47 0.67 2.82 1.04 14.80 a 6768 


























1 Table prepared by Professor S. W. Parr. 


CARBON 281 


The properties of the coals of different kinds follow from their 
composition. Peat, lignites, and bituminous coals increase 
progressively in calorific value as the amounts of moisture and 
oxygen decrease. The oxygen in these coals may be consid- 
ered as combined with either carbon or hydrogen and lessens by 
so much the amount of these elements which can evolve heat by 
combustion. Bituminous coals may equal or even exceed an- 
thracite coals in calorific value because a pound of hydrogen 
gives by its combustion more than three times as much heat as a 
pound of carbon. Such coals, however, give off volatile products 
which burn with a smoky flame, and hence require much greater 
care in use to secure effective combustion. 

Three classes of bituminous coals are distinguished: coking 
coals, which sinter together when heated, giving a hard, coherent 
coke; noncoking coals, which do not sinter, or sinter imperfectly, 
giving a friable coke; and cannel coals, coals of a peculiar, homo- 
geneous structure and conchoidal fracture, which burn with a 
brilliant flame like that of a candle. These last coals are used 
in the manufacture of illuminating gas. The difference between 
coking and noncoking coals seems to be occasioned by the 
presence or absence of some compound whose character is little 
understood and which does not seem to be closely connected with. 
the percentage composition of the coal. 

Chemical Properties of Carbon. The most remarkable prop- 
erty of carbon is the extreme slowness with which it reacts at 
ordinary temperatures with elements for which it has a very 
strong affinity at high temperatures. This is especially true in 
its relation to oxygen. Elementary carbon in either of its three 
forms may remain in contact with air for centuries without any 
apparent effect, although at very high temperatures there seems 
to be almost no element from which carbon will not take away 
oxygen. Practically all organic compounds must be considered 
as in a state of unstable equilibrium in the presence of oxygen, 
for we have only to heat them to their kindling temperature 
when they will burn with very considerable evolution of heat. 
On this property depends the use of carbon and its compounds 


282 A TEXTBOOK OF CHEMISTRY 


for fuel, for the reduction of iron ores and for other metallurgical 
operations. On this property, too, depends the existence of the 
almost infinite variety of compounds which form the material 
basis of the world of life — compounds showing all possible 
gradations in their content of energy and the existence of which 
would be impossible if carbon passed quickly, as most other 
elements do, to the most stable forms of combination. 

In combination with other elements carbon is almost always 
quadrivalent. Methane, CH,, and carbon dioxide, COz, may be 
considered as the most typical compounds. Carbon is bivalent 
in only a very few compounds and exclusively in combination 
with atoms or groups of a negative character, as in carbon 
monoxide, C=O, hydrocyanic acid, H—N=C, and fulminic acid, 
H—O—N=C. There is some evidence that it may be tri- 
valent in very unusual combinations, but it is then extraordi- 
narily reactive. 


CHAPTER XVII 


HYDROCARBONS. ILLUMINATING AND PRODUCER GAS. 
FLAME 


CARBON combines with hydrogen to form many hundreds of 
compounds, called hydrocarbons. These compounds may be 
classified in a number of series in accordance with their com- 
position. The following table illustrates the relations which 
have been found between the formulas of successive hydrocar- 
bons in any series and between the hydrocarbons of different 
series. Each series is named from its first member. In the 
table only one series for a given general formula is given, but, as 
will be seen below, for each general formula, except the first, 
two or more series are possible. The series of the formula 
C,,H>,-4 is omitted from the table because the lower members 
of this series are relatively unimportant. 








Marsg Gas Series | ETHYLENE SERIES ACETYLENE SERIES | BENZENE SERIES 
C C_H CH 


sre, C,, Ho, n~2n—2 n~ 2n—6 














Methane CH, 
Ethane C2H, | Ethene CoH, | Acetylene CoH» — 
Propane C3Hgs | Propene C3H, | Propine C3H, Se 
Butane C.4Hi | Butene C4Hs | Butine C,H —— 
Pentane C;H,. | Pentene C;H,)|Pentine C;Hs ae 
Hexane C,sHy, | Hexene CsHy. | Hexine C,H | Benzene CsHe 
Heptane C,;Hi, | Heptene C;Hy | Heptine Cy7Hi2 | Toluene CH, 
Octane CsHig | Octene CgHy.5 | Octine CsHys | Xvlene CgHio 

















The existence of these compounds may’ be explained very 
simply on the hypothesis that carbon is quadrivalent and that 
carbon atoms unite readily with each other. This hypothesis 
283 


284 A TEXTBOOK OF CHEMISTRY 


gives us the following formulas for the first three members of the 
Marsh gas series : 


H Hive tt Hil aH: 
| aed [ig hgea 
Me ehinekee oi. | 
H A at | ses 
Methane Ethane Propane 


For the fourth member of the series the theory suggests two 
formulas : 


[Pan Raat ele 
H—C—C—C—C—H Hic ioe 
| ipsam ae 
Hes hioe Hest H H 
Faiieel vee +1° Pane 
H 
Isobutane 


Boiling point, —11.5° 


These two hydrocarbons have been prepared by methods 
which leave no doubt as to the structure of each. 

For the series C,H, the theory suggests that we may have 
compounds in which carbon atoms are doubly united and also 
compounds in which there is a ring of carbon atoms. Thus we 
may have: 


H HH ae 
ee C 
| See 
H Fait aN 
Propylene (Propene) Cyclopropane 
Boiling point, —37° Boiling point, —35° 


Both of these compounds are known, and the structure has 
been established by a study of the methods of preparation and 
of their conduct toward various reagents. 


HYDROCARBONS 285 


For the series C,H,» there are four possibilities: one triple 
union, as in acetylene, H—C = C—H; two double unions, as 
in butadiene, CH,—CH—CH=CH:; cyclic compounds with 

CH.—CH 
one double union, as cylopentene, CH, || ; and compounds 
\CH:—CH 


with two cycles, or rings, as dekahydronaphthalene, 


CH,.—CH._CH—_CH.—_-CH, 


| | | 
CH,—CH;—CH—CH,—CH, 


The illustrations given would seem to include all of the types 
of combination possible for carbon and hydrogen atoms, since 
quadruple unions between carbon atoms would be impossible 
for atoms which are united to any other atoms. There are, 
however, certain other relations which seem to depend on the 
arrangement of the atoms in space. 

Some of the combinations of these forms give properties which 
would not be expected from the formulas of the compounds. 
This is especially true of the benzene series, in all of the com- 
pounds of which there is a ring of six carbon atoms, each of which 
is united to one hydrogen atom or to some other univalent atom 
or group. The simplest formula of benzene is that proposed 
by Kekulé: 

H 


C 
cae “us 
HC pe 


| 
H 


and many of the properties and reactions of the hydrocarbon are 
satisfactorily represented by this formula, but for other proper- 
ties it does not give a satisfactory account. A further considera- 
tion of this and similar questions is impossible here. 


286 A TEXTBOOK OF CHEMISTRY 


Marsh Gas or Methane, CH. When decaying leaves in the 
bottom of a pond are stirred, bubbles of a combustible gas con- 
sisting largely of marsh gas or methane, CH,, rise to the surface. 
The same gas escapes from seams of coal, doubtless having been 
formed in a similar manner. In coal mines it is called fire damp. 
A combustible gas consisting very largely of methane is often 
found stored in large quantities in porous sandstones or lime- 
stones lying beneath an impervious layer of shale so situated as 
to form a large inverted reservoir. The gas is usually under 
strong hydrostatic pressure from water beneath. When such a 
reservoir is pierced by boring from above, the gas escapes through 
the opening, and is known as natural gas. When almost any 
kind of organic matter is heated, methane is one of the products 
of decomposition, hence it is always a constituent of illuminat- 
ing gas made by heating coal, oil or wood. At a white heat 
carbon will unite directly with hydrogen to form methane: 


Ct He OH, 


The equilibrium of the reaction is, however, very far on the 
side toward the decomposition of methane into carbon and hy- 
drogen. Water gas (p. 296) usually contains a very small 
amount of methane, which is probably formed by the direct 
union of the elements. 

In the laboratory methane is most easily prepared on a small 
scale by heating a mixture of sodium acetate, NaC2H;O2, and 
soda lime, which is a mixture of sodium hydroxide, NaOH, and 
slaked lime, Ca(OH)>. The slaked lime is added to render the 


mixture infusible: 


NaC.2H302 -- NaOH == Na.CO3 + CH, 
Sodium 
Carbonate 


Methane is the lightest gaseous compound known. It may be 
condensed to a liquid, which boils at — 164°. It is a compara- 
tively stable compound, and its kindling temperature is higher 
than that of hydrogen or than that of most other hydrocarbons. 


SUBSTITUTION 287 


Mixtures of the gas with oxygen or with air explode violently 
when ignited. It burns from a jet with a blue flame, which 
gives very little light. 

Substitution. When a mixture of methane and chlorine is 
exposed to the sunlight, a double decomposition occurs in which 
one atom of the chlorine molecule combines with an atom of 
hydrogen while the other combines with the carbon, apparently 
taking the place of the hydrogen. This process, which occurs 
in a great variety of reactions of organic compounds, is called 
substitution : 


H H 
| | 
Saas +Cl-Cl. = H—C—Cl+ H—Cl 
| 
H H 


Methyl Chloride 


The process may be continued till all of the hydrogen has been 
replaced by chlorine: 
CH;Cl + Cl, = HCI + CHCl, 


Methylene 
Chloride 


CHCl. + Cle = HCl + CHCl; 


Chloroform 


CHCl; + Cly = HCl + CCl, 
Carbon 
Tetrachloride 
Practically, a mixture of the four products is obtained by this 
process, so that these reactions have only a theoretical interest. 
The Davy Safety Lamp. ‘The explosive character of mixtures 
of methane and air has been mentioned. Larly in the nineteenth 
century the frequent explosions of fire damp in coal mines in Eng- 
land, often causing the death of miners, led to a request of Sir Hum- 
phrey Davy that he should investigate the matter and endeavor 
to suggest aremedy. He found, as the result of his investigation, 
that mixtures containing one volume of methane with more than 
six and less than fourteen volumes of air would explode when 


288 A TEXTBOOK OF CHEMISTRY 


ignited.!. Outside of these limits explosions do not so readily 
occur. He also found that the kindling temperature of such 
mixtures is comparatively high, requir- 
ing contact with a surface heated nearly 
or quite to dull redness.2 This prop- 
erty can be easily illustrated for illu- 
minating gas by pressing a cold piece of 
wire gauze down over the flame of a 
Bunsen burner (Fig. 80). Unburned 
gas from the center of the flame will 
pass through the gauze and the mixture 
of gas and air above the gauze will not 
take fire until the latter becomes nearly 
Fig. 80 red-hot. On the basis 
of this fact Sir Hum- 
phrey Davy invented the Davy Safety Lamp, 
which has the flame of the lamp completely 
surrounded by wire gauze (Fig. 81). The 
lamp must, of course, be lighted and closed 
before the miner enters the mine. The lamps 
are usually so constructed that they can be 
locked. Sometimes a lock is used which can 
only be opened with a strong electromagnet, 
so that it will be impossible for the miner to 
open the lamp in the mine. While the danger 
of explosions is greatly lessened, it is not en- > 
tirely removed by the use of the lamp. If Fig. 81 
much fire damp is present, a cap of flame ap- 
pears inside, above the flame of the lamp; and this might, some- 
times, heat the wire gauze to the kindling temperature of the 
mixture. In blasting, too, the sudden vibration from the blast 












ine 


was 


1 Later investigations have shown that these results are only a 
very rough approximation. One authority states that air contain- 
ing 2 per cent of methane may be dangerous. 

2 V. Meyer, many years later, found that the kindling temperature 
of mixtures of methane and oxygen in glass vessels is 650°-680°. 
Ber. 26, 2429. 


PETROLEUM 289 


may carry the flame through so quickly that it is not cooled 
below the kindling temperature by the gauze. 

Mixtures of very fine dust containing organic matter with air 
may explode in the same manner as fire damp. In this way 
destructive explosions have occurred with coal dust in dry mines, 
with flour in flour mills, and with similar dust in other factories. 

Homologues of Methane. On examining the formulas of the 
hydrocarbons given in the table it will be found that the succes- 
sive members of any series differ by one carbon and two hydrogen 
atoms. The reason for this is apparent from the structural 
formulas of methane, ethane, propane, etc. Any hydrocarbon 
in such a series is called a homologue of the lower members of 
the series and the series is called a homologous series. 

Petroleum is found in great underground reservoirs somewhat 
similar to those containing natural gas. Oil fields have been 
found widely distributed in America, especially in Pennsylvania, 
Ohio, Indiana, Illinois, Kansas, Texas, California and Canada. 
A large field is found in the Caucasus, and doubtless very many 
undiscovered fields exist in other parts of the world. Petroleum 
consists chiefly of a very complex mixture of hydrocarbons. 
The petroleum from different localities differs very considerably 
in the nature of the hydrocarbons which it contains and also in 
the amount of the compounds of sulfur which are present. The 
Pennsylvania petroleum consists largely of homologues of me- 
thane. California and Caucasus petroleum contain compounds 
of the cyclic series. 

Crude petroleum is often used as a fuel. Its calorific value is 
about one half greater than that of the best quality of coal. 
Petroleum is refined chiefly by fractional distillation. It is also 
treated with concentrated sulfuric acid to remove compounds 
with a disagreeable odor or objectionable properties. Com- 
pounds of sulfur are removed by boiling it with copper oxide. 
The principal product is usually kerosene, used as a burning 
oil in lamps. The low-boiling products are called petroleum 
ether and ligroin in chemical laboratories, or, commercially, 
gasoline, benzene and naphtha, partly in accordance with the boil- 


290 A TEXTBOOK OF CHEMISTRY 


ing point, but chiefly according to the use made of the material. 
The vapors of the low boiling products form dangerously explo- 
sive mixtures with air. Kerosene should not give enough vapor 
to explode at any temperature below 65° (150° F.), and this is the 
legal flashing point in most states. 

Products boiling at a higher temperature than kerosene are 
used as lubricating oils. Solid products, called paraffin, are 
made into candles and are used for covering jellies and for many 
other purposes. A semisolid product, called vaseline,! is prepared 
for medicinal use. All of the products are very complex mix- 
tures of hydrocarbons. 

Ethylene or Ethene, C.H,. When ordinary alcohol, C,.H;OH, 
and concentrated sulfuric acid are mixed in such proportion that 
the mixture boils at 140°, ethyl ether, (C2Hs).0, and water distill 
over on heating. If more sulfuric acid is used (6 parts of sulfuric 
acid to 1 of aleohol by weight) so that the mixture distills or de- 


H Va: 
composes at 170° — 180°, ethylene, Se = ~ ,isformed. In 
H H 


both cases we may consider that the sulfuric acid removes water 
from the alcohol, but the mechanism of the reaction is more com- 
plicated than such a statement indicates. 


2 C.H;OH ae H.O = C.H;—O—C 2H; 
Ethyl Ether 


C.H;0H ee H,O == Geli 
Ethylene 


Ethylene is a colorless gas, very slightly lighter than air. It 
has a sweetish odor and burns with a bright, luminous flame. 
The difference between methane and ethylene in this regard 
seems to depend on the fact that methane is quite stable, even at 
comparatively high temperatures, and does not readily decom- 
pose with the separation of carbon, while at the temperature of 
the flame ethylene decomposes, partly, into methane and carbon: 

1 Vaseline is a proprietary name used by the Chesebrough Manu- 


facturing Company. The name used in the Pharmacopocea and 
by other manufacturers is petrolatum. 


UNSATURATED COMPOUNDS 291 
OaHy = CH, -f C 


The carbon which is liberated temporarily assumes the solid 
form and, being raised to a white heat by the flame, makes 
it luminous. In spite of this instability, ethylene is formed 
when any hydrocarbon of the methane series is heated to a 
high temperature or, indeed, when almost any organic com- 
pound is heated. [or this reason it is always present in 
illuminating gas prepared by heating coal or oil and is one of 
the most important constituents of the gas, because of the lumi- 
nous quality of its flame. It is formed to some extent even 
from methane, although it is less stable than methane and the 
reaction : 


2 CH, = CoH, + 2 H, 


isendothermic. It would almost seem that this ability of carbon 
to enter into many reactions inewhich heat is absorbed from sur- 
rounding objects is one of the most important characteristics 
of the element. It is doubtless intimately connected with those 
properties of the carbon atom which cause the rate of many of its 
reactions to be so slow (p. 281). The result of this seems to be 
that the speed of a reaction in one direction or another often has 
more effect in determining the direction of the reaction than the 
heat evolved or absorbed. | 

Unsaturated Compounds. Ethylene Chloride and Ethylene 
Bromide. Ethylene combines directly with chlorine to form 
ethylene chloride, C2H1Cle, and with bromine to form ethylene 
bromide, C2H.Brye. The process is known as addition. These 
compounds may be called, also, dichloroethane and dibromo- 
ethane and are to be considered as substitution products of 
ethane, C.Hs. They illustrate the tendency of compounds 
having double or triple unions between carbon atoms to take 
up other elements and pass back into compounds which are 
derivatives of the hydrocarbons of the methane series. For 
this reason the hydrocarbons of the ethylene and acetylene series 
are called wnsaturated, while the hydrocarbons of the methane 


292 A TEXTBOOK CF CHEMISTRY 


series and their derivatives are called saturated. This conduct 
ist el 
| 


has led some chemists to prefer the formula H—C—-C—H for 
H H 


ethylene instead of the usual formula, dent ele Which- 
ever formula is true, it is evident that carbon atoms which are 
spoken of as doubly united are not more firmly held together 
than by a single union. The reverse of this seems to be true. 
The double union is a point of especial reactivity. 

Acetylene. When an electric arc is formed between carbon 
points in an atmosphere of hydrogen, some acetylene, CoHa, 
is formed. Acetylene is an endothermic compound and de- 
composes into carbon and hydrogen with evolution of heat: 


CoH, = 2C + He + 53,000 small calories 


for 26 grams of acetylene. The formation of acetylene is often 
used as an illustration of the fact that a high temperature is 
favorable to the formation of endothermic compounds. 

Acetylene is formed by the incomplete combustion of ethylene 
or of carbon compounds generally and so is found among the 
gases coming from a Bunsen burner burning at the base. The 
unpleasant odor of these gases is not, however, due to the acety- 
lene. The presence of the acetylene can be shown by driving 
the gases through an ammoniacal solution of cuprous chloride, 
with which the acetylene gives a precipitate of copper carbide, 
often incorrectly called copper acetylide: 


CucsCl. + ¥ NH; -f CH. = CusCe + 2 NH,Cl 
Copper 
Carbide 


The formation of copper carbide in this way indicates that 
acetylene has some of the properties of an acid. ‘The electri- 
cal conductivity of solutions of acetylene in water also indicates 
that it is an acid, but an extremely weak one, so that its salts 


ACETYLENE 293 


are hydrolyzed by water. On this fact depends its prepara- 
tion for commercial uses from calcium carbide and water: 


CaC, — 2 HOH = Ca(OH). 4- CH. 
Calcium 
Carbide 


Acetylene is a colorless and odorless gas, which may be con- 
densed to a liquid or solid by cold or pressure. The boiling 
point is — 83.6° and the melting point a little higher, — 81.5°. 
Under atmospheric pressure, therefore, it sublimes without 
melting. | 

Acetylene is unsaturated in the same sense as ethylene and 
may combine with four atoms of bromine to form acetylene 
tetrabromide or tetrabromoethane, C2H2Bra. 

Acetylene dissolves readily in acetone, one volume of the 
liquid dissolving 25 volumes of the gas at 15° under atmos- 
pheric pressure and 300 volumes under a pressure of 12 atmos- 
pheres. As a solution containing 100 volumes of the gas for 
one of acetone is not explosive (see below) such a solution can 
be used to advantage for illuminating purposes. 

Acetylene burns from an ordinary jet with a smoky flame, 
due to the separation of carbon. When burned from a suitable 
burner it gives a very brilliant white light, developing from 
twelve to fifteen times as much light as can be obtained with 
the same volume of good illuminating gas burned with an ordi- 
nary burner. The intense light is due partly to the ease with 
which acetylene decomposes into carbon and hydrogen, but 
doubtless also to the heat developed by the decomposition, 
which aids in raising the temperature of the particles of carbon 
to a white heat. It is estimated that while only about 2 per 
cent of the energy of illuminating gas is actually effective as 
light, about 10 per cent of the energy of burning acetylene may 
appear as light. 

At 20° acetylene may be condensed to a liquid under a pres- 
sure of 41 atmospheres. This is considerably less than the 
vapor pressure of liquid carbon dioxide at the same tempera- 


294 A TEXTBOOK OF CHEMISTRY 


ture, and when a process had been invented by means of which 
the manufacture of calcium carbide in an electric furnace be- 
came commercially possible it was thought that liquid acetylene 
could be very conveniently used in the liquid form, condensed in 
strong steel cylinders. Soon after the first attempts in this direc- 
tion were made, however, some unexpected and, at first, unac- 
countable explosions occurred. An investigation of the matter 
soon showed that liquid acetylene, or, indeed, gaseous acetylene, 
even under a pressure of somewhat less than two atmospheres, 
may be exploded by a glowing wire or by a fulminating cap. 
The explosion is due, of course, to the fact that acetylene de- 
composes into carbon and hydrogen with considerable evolution 


of heat: Cre oc eT 


Although the volume of the hydrogen is the same as that of 
the acetylene the heat of decomposition increases the pressure 
and is sufficient to cause the decomposition, when once started, 
to proceed explosively from one part of the liquid or compressed 
gas to another. It has been pointed out above that solutions 
of acetylene in acetone are not explosive, if the concentration 
is not carried too far. Such solutions are now extensively used, 
especially for automobile and motorcycle lights. 

At slightly elevated temperatures acetylene polymerizes 
easily, that is, it combines with itself to form more complex 
compounds, some of which are liquid or solid at ordinary tem- 
peratures. When acetylene is generated by dropping water 
on calcium carbide the heat evolved by the reaction may cause 
a very considerable loss by polymerization. For this reason 
those forms of generators in which the carbide is dropped inte 
the water are most suitable. 

Benzene, CsHs. In the manufacture of illuminating gas by 
the distillation of bituminous coal and also in the manufacture 
of coke by the methods in which retorts are used and the by- 
products are saved, large quantities of coal tar are produced. 
This is an extremely complex mixture from which many valuable 
products are obtained, chiefly by fractional distillation. Among 


ILLUMINATING GAS 295 


these products are benzene and its homologues and some other 
hydrocarbons, especially naphthalene and anthracene. These 
hydrocarbons are very extensively used in the manufacture of 
the coal-tar dyes and of many other compounds which are used 
in medicine and in the arts. Benzene may also be formed by 
the polymerization of acetylene: 


3 CH. = C.He 


It is a colorless liquid which melts at 5.4° and boils at 80.2°. 

Illuminating Gas. The manufacture of illuminating gas by 
heating bituminous coal, especially cannel coal, in earthen- 
ware or iron retorts, has been repeatedly referred to. The gases 
obtained in this manner consist chiefly of a mixture of hydrogen, 
methane, carbon monoxide, carbon dioxide and hydrogen 
sulfide, with a small per cent of the so-called “‘ heavy hydro- 
earbons.”’ These last consist of ethylene, C2H4, vapor of 
benzene, CeHe, a little acetylene, C2Hs, and small amounts of 
many other gases and vapors. In burning the gas only the 
heavy hydrocarbons decompose to an appreciable extent with 
separation of carbon. On this account these hydrocarbons are 
sometimes designated as illuminants. 

The composition of the gas varies greatly both with the tem- 
perature and the duration of the heating of the coal. A high 
temperature favors the decomposition of the hydrocarbons and 
increases the per cent of hydrogen, reducing the illuminating 
power of the gas, but it also greatly increases the volume of the 
gas obtained from a given weight of coal. The gas which es- 
capes from the coal during the first part of the heating is also 
much richer in the heavy hydrocarbons than that given later. 
These facts are easily understood from the instability of the 
hydrocarbons, at high temperatures. 

The hydrogen sulfide in the gas must be removed by passing 
it through boxes or chambers having a series of shelves covered 
with slaked lime, Ca(OH)., or more often by passing it through 
boxes containing moist ferric hydroxide, Fe(OH); : 


2 Fe(OH); + 3HS = 2FeS + S+3H,0 


296 A TEXTBOOK OF CHEMISTRY 


The ferrous sulfide passes back into a mixture of ferric hydrox- 
ide and sulfur on exposure to the air and so may be used re- 


peatedly: 9 eS + 30+ 3H,0 = 2 Fe(OH); +28 


The illuminating power of the gas is determined by com- 
parison with the light of a standard spermaceti candle which 
burns 120 grains per hour. For comparison, the gas is burned 
at the rate of 5 cubic feet per hour. Gas of good quality should 
give from 18 to 22 candle power. 

The light which can be obtained from a given quantity of 
gas may be greatly increased by burning it from a Bunsen burner 
under a mantle composed of oxides of thorium and cerium — 
the Welsbach light, named after the inventor. The oxides not 
only furnish the solid substance heated to a high temperature, 
which is necessary in almost all forms of practical illumination, 
but they also catalyze the reaction of combustion, localizing 
the latter in immediate contact with the solid particles and so 
greatly increasing the temperature to which these are raised. 
(See p. 364.) With the inverted Welsbach burner, gas may give 
more than ten times as much light as could be obtained with an 
ordinary flat flame. 

Oil Gas. When petroleum is heated to a high temperature 
the hydrocarbons which it contains are decomposed with the 
formation of carbon, hydrogen, methane and other hydrocar- 
bons, some of which are gaseous at ordinary temperatures. If 
the temperature is high enough and especially with the aid of 
certain catalyzers, the final products are hydrogen and carbon. 
The process has been proposed as a method to obtain hydrogen 
for filling balloons. At lower temperatures it is possible to 
obtain a gas very rich in the heavy hydrocarbons, and a gas, 
called Pintsch-gas, is manufactured in this way and compressed 
in steel cylinders for use in lighting railway coaches and for 
similar purposes. 

Water Gas. When steam is passed over incandescent coal 
or coke a mixture of carbon monoxide, CO, and hydrogen, He, 


is formed: C ate H.O — GO ve H, 


PRODUCER GAS 297 


The heat of combustion of the carbon monoxide and hydrogen 
is very much greater than the heat of combustion of the carbon 
or coke. In other words the reaction is endothermic in a very 
high degree. The mass of incandescent coke cools very rapidly 
as the reaction proceeds. Practically, the reaction is carried 
out intermittently. The coke, contained in a large chamber, 
is brought to a white heat by burning a part of it in a blast of 
air, while the products of combustion are allowed to escape or 
are utilized as a fuel gas. The blast of air is then shut off, 
steam is turned on and the mixture of carbon monoxide and 
hydrogen, called “water gas” is collected for use. After a 
few minutes the mass cools below the temperature of rapid 
reaction. The steam is then shut off and the heating process 
repeated. By this method a gas having about one half of the 
heating value of a good illuminating gas can be manufactured 
very rapidly and cheaply. It is not suitable for use as an il- 
luminating gas, since hydrogen burns with a colorless flame 
and carbon monoxide with a blue flame which gives very little 
light. It may be enriched, however, by the addition of oil gas 
and in that form is used as illuminating gas in many cities of 
the United States. The most serious objection to its use is 
the very poisonous character of the carbon monoxide which 
it contains. Not only is this dangerous because of the acci- 
dental escape of gas from an open stopcock, but in the winter 
time gas may escape from leaks in pipes and may travel for 
some distance beneath the frozen ground till it finds an outlet 
in a cellar. The odor characteristic of illuminating gas is re- 
moved by passage through earth, still further increasing the 
danger. For these reasons some states forbid the manufac- 
ture of a gas containing more than a stated, small per cent of 
carbon monoxide. 

Producer Gas. ‘The water-gas process is a wasteful one from 
the point of view of the per cent of the energy of the coke finally 
obtained in the water gas, largely because of the thick layer 
of fuel which must be used. Under such conditions the carbon 
is burned during the heating stage only to carbon monoxide, 


298 A TEXTBOOK OF CHEMISTRY 


CO, the carbon dioxide which is formed in the lower part of the 
chamber being reduced to carbon monoxide above: 


CO, + C = 2CO 
The heats of combustion involved are as follows : 


Amorphous C + Og = COz + 97,650 small calories 
CO + O = CO, + 68,200 small callories 
Hence, amorphous C + O = CO + 29,450 small calories. 


It is evident from this that less than one third of the heat 
energy of coke is utilized when it is burned only to carbon 
monoxide. These heat relations, which are so unfavorable 
to the economy of the water-gas process may be utilized for 
the production of a low-grade fuel gas, commonly called “ pro- 
ducer gas,’ which may, under favorable circumstances, retain 
from 80 to 85 per cent of the original heat energy of the coal 
or fuel employed in its manufacture. A chamber containing 
a thick bed of fuel has a blast of moist air forced through it in 
such a manner that a gas consisting chiefly of carbon monoxide 
and nitrogen, with a little hydrogen is obtained. The heat 
energy of such a gas can be much more perfectly utilized than 
that of solid fuel for many metallurgical operations, for the melt- 
ing of glass, for use in specially constructed gas engines and for 
many other purposes. 

Blast-furnace Gas. The reduction of oxides of iron by hydro- 
gen has been spoken of as a reversible reaction. The same is 
true when carbon monoxide is the reducing agent : 


FeO; aa 3CO@2Fe + 3 CO. 


Whichever reaction is used, the equilibrium of the reaction 
is so far to the left that the process can be successful only in 
the presence of a very large excess of hydrogen or carbon monox- 
ide. For this reason the gases escaping from the top of a 
blast furnace (p. 541) are of much the same nature as producer 
gas, with the advantage that the oxygen of the carbon monox- 
ide which they contain comes partly from the iron ore and so 


ILLUMINATING AND PRODUCER GAS 299 


che per cent of nitrogen may be lower. This gas has long been 
utilized for heating the blast, generating steam, etc. During 
recent years it is coming into extensive use in gas engines. 

The following table illustrates the composition of the various 
kinds of gas which have been mentioned in this chapter: 
































GoiL On ENRICHED] PRo- Buast 
Can ae WaTER | DUCER | FURNACE 
Gas Gas Gas 
Carbon dioxide, CO. . . .| 1.1 | —— 3.0 Lepage 
Carbon monoxide, CO. . .) 7.2 | —— 26.1 Dot 28.0 
Hydrogen, H, . 7. . . .| 49.0 | (14:6]") 32.1 6.0 Def 
Prerianier Ola vscry ho he” i 8425 B88 19.8 3: 0.2 
“Heavy Hydrocarbons” . .| 5.0 | 45.5 16.6 aan sa 
Nitrogen ee Che yo Ds Lease oy pect 66.0 | 57.1 
Candle Power... . . | 17.5 | 65.0 | 25.0 








1 Hthane, CoH. 


Luminous Flames. It was formerly supposed that carbon 
separates in a flame because the hydrogen is more easily burned 
than the carbon. A study of the equilibrium between the gases 
present in a flame has shown, however, that this is not the case. 
It is only the carbon which results from the dissociation of 
hydrocarbons and which is momentarily heated to a very high 
temperature, which gives the luminous quality to the flame. 
In a gas flame burning from a round opening we may distinguish 
.clearly three parts: 

1. The interior cool portion of unburned gas. The head of a 
match inserted to this point enflames slowly or not at all. 

2. A mantle of partial combustion and of dissociation of 
hydrocarbons, the luminous part of the flame. 

3. A blue mantle surrounding the whole and especially notice- 
able at the base, where complete combustion to carbon dioxide 
and water occurs. 

The same parts may be seen in a candle flame, the gas or va- 
pors in the interior being formed by the heat of the flame acting 
on the materials of the candle. 


300 A TEXTBOOK OF CHEMISTRY 


Bunsen Burner. A luminous flame may be made nonlumi- 
nous or slightly luminous by either of two methods. The flame 
may be diluted by an indifferent gas till it is cooled below the 
temperature of rapid dissociation for the hydrocarbons present, 
or oxygen may be introduced in such quantity that the carbon 
burns at once to carbon monoxide. Both effects are present 
in the Bunsen burner. The introduction of air at the base of 
the burner causes the combustion to take place in two stages. 
In the inner cone, Fig. 84, p. 303, the constituents of the gas are 
partially burned, giving a mixture of nitrogen, hydrogen, car- 
bon monoxide, carbon dioxide and water vapor. The last four 
substances are present above this cone in accordance with the 
equilibrium of the reversible reaction : 


CO + H.O = CO; + H: 


At a temperature of about 850°-900° the four compounds 
will be present in equal amounts, by volume, if the three ele- 
ments are present in the proportion given in the equation. 
Stated otherwise, at this temperature the reducing power of 
carbon monoxide is equal to that of hydrogen and the oxidiz- 
ing power of water vapor is the same as that of carbon dioxide. 
At lower temperatures the equilibrium is displaced to the right 
because the carbon monoxide becomes the stronger reducing 
agent, while at higher temperatures the equilibrium is displaced 
to the left. It is in accordance with this equilibrium that 
water gas prepared at a high temperature contains very little: 
earbon dioxide. This is of considerable importance for an 
illuminating gas, because a small per cent of carbon dioxide 
greatly reduces the illuminating power. At the outer mantle 
of the Bunsen flame the carbon monoxide and hydrogen burn 
to carbon dioxide and water. By the arrangement shown in 
Fig. 82 it is possible to separate the two zones of the Bunsen 
flame from each other in such a way that the gases between the 
two zones may be drawn out and analyzed. By this method it 
has been shown that the equilibrium agrees with that of the 
water-gas reaction given above. 


ILLUMINATING AND PRODUCER GAS 301 


In carrying out the experiment it is well to make the tube of 
the Bunsen burner, shown at the base of the figure, about 
60 cm. long, to secure thorough mixture of the gas and air, and 
the proportion between the two must be properly regulated. 
If a spray of a solution containing a lithium and a copper salt 
is introduced into the gas, the lower cone will be colored 
red from the lithium and the upper cone green from 
the copper, because lithium is oxidized by even the 
small amount of oxygen in the water, carbon monoxide 
and carbon dioxide of the lower cone, while the copper 
is not oxidized till the freer oxygen supply of the upper 
‘cone is reached. Smithells and Ingle, J. Chem. Soc. 
61, 204 (1902) ; Smithells, Phil. Mag. [5] 39, 132 (1895). 

When the proportion of air entering a Bunsen burner 
at the base is increased, the inner cone grows shorter 
until a point is reached where the flame “ snaps back ” 
and burns at the base. This is because the inner cone 
is to be looked upon as a stationary explosion wave 
where the velocity of the current of gases upwards is 
equal to the velocity of the combustion downwards. 
An increase in the proportion of oxygen increases this 
combustion velocity till it exceeds the velocity of the 
current of gas. 

Explosion Waves. When a mixture of gases is ex- 
ploded the combustion is not, of course, instantaneous, 
but proceeds with a definite velocity from the point of 
ignition. This phenomenon has been carefully studied, 
partly by measurements of the velocity by stationary flames, as 
suggested above (Michelson, Wied. Ann. 37, 19 (1889)), partly 
by measuring the pressures developed (Michael and Le Chatelier, 
Ann. des Mines [8] 4, 379, 599 (1883)), partly by photographing 
the flame by its own light (Dixon, Ber. 38, 2426 (1905), Phil. 
Trans. 184, 97 (1893); 200, 315 (1903)). It has been found 
that there are two distinct stages in the explosion. At first the 
combustion proceeds with only a moderate velocity —for a 
mixture of carbon monoxide and oxygen from 20 to 91 cm. per 





Fig. 82 


302 A TEXTBOOK OF CHEMISTRY 


second, according to the composition of the mixture. The ex- 
pansion of the burning mixture by the heat compresses the gas 
in front of the explosion wave. This adiabatic! compression 
raises the temperature of the gas and, according to the prin- 
ciples of thermodynamics, in cases where the final volume is 
diminished by the explosion, lowers the explosion temperature. 
If the apparatus in which the explosion occurs is large enough, 
a point will finally be reached at which the compression raises 
the temperature of the mixture in front, to its explosion temper- 
ature. When this point is reached, the velocity suddenly 
changes and the flame proceeds with the velocity of sound. 
The pressure of the compression wave may become ten times 
that of the original mixture and it is estimated that temperatures 
of 5000°-6000° may be produced. The velocity of sound at 
5000° would be about 1400 meters per second and the velocities 
actually measured somewhat exceed this. This result is of 
considerable importance in connection with the design of gas or 
gasoline engines, where the form and dimensions of the explosion 
cylinder should be chosen so as to avoid the severe shock which 
comes with the second type of explosion waves. it also explains 
the shattering of a long eudiometer at the end farthest from the 
point ef ignition. A glass tube of moderate thickness will readily 
withstand the explosion of a mixture of oxygen and hydrogen, 
if the dimensions are such as to avoid the explosion wave of the 
second type. 

Temperature of Flames. The temperature of the flame of a 
Bunsen burner varies from about 300° in the center, near the 
mouth of the burner, where combustion has not begun, to about 
1550° in the portion between the inner cone and the outside of 
the flame. These temperatures are shown in detail in Fig. 84. 

In the Meker burner, Fig. 83, by widening the top of the 
burner and giving it a considerable number of fairly heavy 
metallic partitions, the inner cone is divided into a number of 


1 Adiabatic means without escape of heat. Here the compres- 
sion is so rapid that the heat which results from the compression 
cannot escape. 


ILLUMINATING AND PRODUCER GAS 303 


small and very short parts. This brings the high temperature 
of the upper part of the Bunsen flame down close to the mouth 
of the burner, concentrates the flame and gives it a more uni- 
form and somewhat higher temperature. 

The temperatures given in the figures are, of course, the tem- 
peratures of the flame when no substance radiating heat is 




















Fig. 83 Fig. 84 


present. A platinum or porcelain crucible placed in the flame 
will be at a much lower temperature. A 20-gram platinum 
crucible placed 1 cm. above the Meker burner, with ordinary 
gas, will usually have a temperature of 900°-950°. 

Blowpipe. By means of a blowpipe (Fig. 85) the flame of a 
candle or of a Bunsen burner may be conveniently used for 


304 A TEXTBOOK OF CHEMISTRY 


heating or for oxidizing or reducing substances supported on 
charcoal or contained in beads of borax or of sodium meta- 
phosphate (from microcosmic salt). The interior of the flame, 
especially if it retains a slightly luminous character, will have a 
strong reducing effect, while at a point just beyond the tip of the 
flame, where substances are 
heated and at the same time 
can receive oxygen from the 
air, an oxidizing effect will 
be produced. These effects 
ean be readily shown with 
litharge, PbO, and metallic lead. Similar effects can be ob- 
tained in the Bunsen flame. The tip of the inner cone is 
reducing, while the outer edge of the flame is oxidizing. 
Reversed Flames. Under ordinary conditions a flame con- 
sists of a combustible gas surrounded by air or oxygen with 
which it is combining. Because 
the oxygen of the air is available 
without expense further than the 
apparatus necessary to utilize it, 
we have become accustomed to 
speak and think of the combust- 
ible gas as the source of the energy 
which we use. We can, however, 
conceive of a world where the at- 
mosphere should consist of meth- 
ane or some other gas or gases 
which we call combustible. In 
such a universe we might obtain 
energy by preparing and burning 
oxygen. Such a condition can ae Fig. 86 
be illustrated with the apparatus 
shown in Fig. 86. When the illuminating gas entering through 
one tube is in excess, oxygen entering through the other will, if 
ignited, burn in the atmosphere of illuminating gas, while the ex- 
cess of the gas will burn at the end of the lamp chimney above. 


Fig. 85 





Oxygen 


ILLUMINATING AND PRODUCER GAS 305 


EXERCISES 


1. In what proportion must oxygen be mixed with the following 
gases, or vapors considered as gases, for their complete combustion ? 
Methane, ethane, ethylene, acetylene, butane, benzene, gasoline if it 
has the average composition of heptane. 


2. In what proportion must the same gases or vapors be mixed with 
air for their combustion ? 


3. The heats of combustion of hydrogen, amorphous carbon and of 
some carbon compounds are : 

Hydrogen, H; + 302 = H.O + 68,400 small calories. 

Amorphous carbon, C + O2 = COz + 97,650 small calories. 

Methane, CH, + 20, = CO, + 2 H.O + 214,000 small calories. 

Ethylene, C2.H, + 302 = 2 CO, + 2 H2O + 333,350 small calories. 

Acetylene, CoH, + 24 O2 = 2 CO + HO + 313,800 small calories. 

Benzene (vapor), CsHs + 743 O2 = 6 CO2 + 3 H20 + 799,350 small 
calories. 

These values are all for a final condition of liquid water at 18°. 

If the following reactions could occur at ordinary temperatures, how 
much heat would be evolved or absorbed by each ? 


C+2H. = CH, 
CoH, = G + CH, 
CoH. == iplad, + He 
CoH, = 2C +2 He 
3 CH; = CoH 
C + H.0 = CO + He 
C +2H.0 = CO, + 2 He 


CHAPTER XVIII 


OXIDES AND SULFIDES OF CARBON. ASSIMILATION 
AND RESPIRATION. CYANIDES. 


Carbon Dioxide. The sources of carbon dioxide in the air 
and the maintenance of a small constant amount, which fur- 
nishes carbon for the growth of plants, have been referred to in 
a previous chapter. For laboratory uses carbon dioxide is pre- 
pared by the decomposition of carbonates. Two properties of 
carbonic acid lead to the decomposition of carbonates when 
treated with almost any of the common acids. It is an exceed- 
ingly weak acid, the ionization : 


1.CO; = Ht HCO; 


taking place to only a very trifling degree, even in dilute solu- 
tions, and it is also very unstable, dissociating so readily to 
carbon dioxide and water: 


yon 
O=C = O=C=O0 + H—O—H 
Noa 

that the acid can exist only in solution, or possibly at very low 
temperatures. 

The carbonates usually employed in the laboratory are cal- 
clum carbonate, CaCQ3, and sodium bicarbonate, NaHCQs3. 

Carbon dioxide is also readily prepared by heating sodium 
bicarbonate, NaHCOs;, magnesium carbonate, MgCOs, or cal- 
cium carbonate, CaCO3. The first dissociates to normal sodium 
carbonate, NaeCOs3, carbon dioxide and water at a compara- 
tively low temperature. Magnesium carbonate requires a some- » 
what higher temperature and calcium carbonate must be heated 
to 812°, or very bright redness, before the dissociation pressure 


of the carbon dioxide is equal to atmospheric pressure. 
306 


CARBON DIOXIDE 307 


Carbon dioxide is a colorless gas with a slightly sour taste 
and odor. It may be condensed to a liquid by pressure or to 
a solid by cold. The solid melts at — 56.4° at a pressure of 5.1 
atmospheres, while its vapor pressure is one atmosphere at —79°. 
The gas can be liquefied, therefore, only under pressures greater 
than five atmospheres. If liquid carbon dioxide is allowed to 
escape from a cylinder in which it is kept under pressure, the 
evaporation of a portion of the liquid will cool the remainder 
below its freezing point and the solid carbon dioxide, which has 
a temperature of —79°, can be collected in a small sack of 
closely woven cloth placed over the nozzle of the cylinder. 
Mixtures of the solid with alcohol, ether or acetone are very 
effective for securing low temperatures. The mixture with 
acetone boils at —88° and a temperature of —110° can be ob- 
tained by forcing air through it. 

Isothermals of Carbon Dioxide. The critical temperature of 
carbon dioxide is 31.5°. In the accompanying diagram, Fig. 87, 











a] 
oS 


jer) 
j=) 










or 
j=) 


i 
=) 








PRESSURE IN ATMOSPHERES. 
oe) 
j=) 


bo 
© 








a 
=) 


VOLUMES. 
Fig. 87 


308 A TEXTBOOK OF CHEMISTRY 


pressures are represented by the ordinates, volumes by the 
abscissas and temperatures by isothermal lines, which show the 
relation between pressure and volume for a given quantity of 
carbon dioxide. Such a diagram shows clearly all of the most 
important relations between pressure, volume and temperature 
for a typical gas. In the region of A where the pressures are 
moderate and the temperatures are very considerably above 
the temperature of liquefaction, for these pressures, the volumes 
are very nearly inversely proportional to the pressures, accord- 
ing to the law of Boyle, and the distance between the isother- 
mals either parallel to the axis of temperatures or parallel to 
the axis of pressures is closely proportional to the absolute tem- 
perature in accordance with the law of Charles. The higher 
the temperature and the lower the pressure, the more nearly 
are these laws accurate. In the region B there are two phases, 
liquid and vapor, and as the total volume is independent of the 
pressure, the isothermals are parallel to the axis of volumes. 
From this point the region where the gas laws are valid is 
reached either by an increase in temperature or a decrease in 
pressure. The latter can occur only when the liquid phase has 
disappeared. In the region C the isothermals are above the 
critical temperature and, while there is a flattening of the curves, 
showing some tendency to liquefy, there is no part parallel to 
the axis of volumes. In the region D the volume is less than 
that occupied by the liquid even at a considerably lower tem- 
perature, but it cannot be said that the substance is liquid, 
since no change of pressure, while the temperature is constant, 
will cause the separation into two phases. The coefficient of 
compressibility, however, approaches that of liquids. 

Density of Carbon Dioxide. The weight of a gram molecular 
volume of carbon dioxide is about 44, while that of air is about 
28.9. Hence the gas is a little more than one half heavier than 
air. When the gas escapes somewhat rapidly from crevices in 
the earth, as is sometimes the case in wells and caves or mines, 
it may accumulate in sufficient amount to suffocate men or 
animals. In the Grotto del Cano in Italy the gas accumulates 


CARBONIC ACID 309 


near the floor of the cave in such a manner that a dog entering 
the cave is suffocated, while a man with his head higher up may 
escape injury. The presence of a suffocating mixture in a well 
can be detected by lowering a candle into it, though a candle 
may be extinguished where respiration for some time is still 
possible. 

The conduct of carbon dioxide in wells and caves has given a 
popular impression that the gas will accumulate near the floor of 
a poorly ventilated room, but owing to the rapid diffusion of 
gases and because the conditions of breathing and flames cause 
an immediate mixture with the air, no appreciable accumulation 
of this sort can occur. 

Aqueous Solutions of Carbon Dioxide, Carbonic Acid. At 
ordinary temperatures water dissolves approximately its own 
volume of carbon dioxide (1.02 volume at 15°, 0.88 volume at 
20°). In accordance with Henry’s law the weight of the gas 
absorbed is very nearly proportional to the pressure and since, 
by the law of Boyle, the weight of a given volume of a gas is 
proportional to the pressure, water will absorb its own volume 
of the gas either at low or high pressures. By saturating water 
with the gas under high pressures a liquid is obtained from which 
carbon dioxide escapes with effervescence on relieving the 
pressure. Some mineral waters of this kind, especially Apolli-, 
naris water, and Congress water from a spring in Saratoga, are 
found in nature, and similar waters are prepared artificially 
either for use directly or as the basis of the so-called “ soda 
water.” The carbonic acid, H2CO;3, formed by the union of 
the carbon dioxide with the water imparts to it a sightly sour 
taste. It also gives an acid reaction, which may be shown by 
the reddening of litmus or the discharge of the color of a faintly 
alkaline solution of phenol phthalein. The carbon dioxide is 
expelled by boiling and the acid reaction disappears. 

Carbonates and Bicarbonates. Hard Waters. The acid 
character of a solution of carbon dioxide is also shown by the 
formation of carbonates or bicarbonates (acid carbonates) with 
bases ; 


310 A TEXTBOOK OF CHEMISTRY 


2 NaOH + H.CO3 = NaeCO3 + 2 HOH 
Sodium 
Carbonate 


Ca(OH). + H.CO; = CaCO; + 2 HOH 
NaOH + H2CO3 = NaHCO; + HOH 


Sodium 
Bicarbonate 


The ionization of the bicarbonate ion: 
HCO;; 22 H* +: CO;> 


is so very slight that very few carbonate ions can exist in aqueous 
solutions. Normal carbonates are therefore hydrolyzed by 
water, and their solutions have an alkaline reaction: 


NazCO; = 2Nat + CO; 
CO; + HOH = HCO,- + OH- 
or | NasCO; + HOH = Nat+ Nat + HCO;- + OH- 


Bicarbonates may also be formed by the action of carbonic 
acid on carbonates : 


CaCO; + H.COs3 — Ca(HCOs). 


The carbonates of calcium and magnesium are only very 
slightly soluble in water (CaCQOs in 77,000 parts), but the acid 
carbonates, or bicarbonates, are much more easily soluble. 
Natural waters, which always contain carbonic acid, partly 
absorbed from the air but chiefly formed by the oxidation of the 
organic matter of the soils under the influence of bacteria, take 
up calcium carbonate and magnesium carbonate from the soils 
in the form of bicarbonates. Such waters are called “ hard” 
waters. When boiled, owing to the ease with which the bicar- 
bonates dissociate into carbon dioxide, water and the normal 
carbonate, such waters give a deposit of calcium and magnesium 
carbonates, which forms the scale in teakettles and steam 
boilers. When the hardness is due to bicarbonate alone, it is 
called “ temporary hardness’ and can be largely removed by 
boiling the. water : 


Ca(HCQs)> = CaCO; + H.O + CO, 


CARBON MONOXIDE 311 


Calcium sulfate, CaSO., which is also present in many natural 
waters, is precipitated only when the water is concentrated or 
when it is heated to a high temperature under pressure. A water 
containing calcium sulfate is said to be “ permanently hard.” 
The precipitation at high temperatures is due to the decreased 
solubility of calcium sulfate under these conditions, and the 
scale formed is particularly adherent and objectionable. Such 
waters may be softened by the addition of an alkaline carbonate, 
phosphate, fluoride or borate, any one of which will precipitate 
the calcium as a nearly insoluble compound : 


CaSO, + NaesCO3 = NaeSO,4 + CaCO; 
3 CaSO4 + 2 Na3PO4 = Ca3(PO.u)2 + 3 NaoSO, 
CaSO, + 2 NaF = CaF, + Na.SOx, 
CaSO, + Na2B.O; = CaBsO, + NaSO, 


Carbon Monoxide. The formation of carbon monoxide in 
the manufacture of water gas and of fuel gas has already been 
‘discussed. The most familiar occurrence of the gas is probably 
in the burning of hard coal, where the carbon monoxide, formed 
by the reduction of carbon dioxide in the interior of the mass 
of coal, burns at the surface with a blue flame. 

Carbon monoxide is most easily prepared in the laboratory 
by the decomposition of oxalic acid, H2C2O.. Concentrated 
sulfuric acid assists in the decomposition, as a dehydrating agent : 


H.C2O,4 = H.O +- CO, oe CO 


The carbon dioxide must be removed from the mixture by 
means of soda lime or by passing the gases through a wash bottle 
containing of sodium or potassium hydroxide. 

Carbon monoxide is a colorless and odorless gas, which may be 
condensed to a liquid that boils at — 190° and toa solid, which 
melts at — 199°. Itburns with a characteristic blue flame, and 
mixtures of it with air or oxygen explode violently. Curiously 
enough, however, a perfectly dry mixture of carbon monoxide 
and oxygen will not explode, and the gas dried with phosphorus 
pentoxide will not burn in air or in oxygen which has been dried 


312 A TEXTBOOK OF CHEMISTRY 


with the same agent. This recalls the fact that dry chlorine 
does not act on iron or copper. Many similar facts for which 
there is at present no very satisfactory explanation are known. 
The effect of a minute trace of moisture in promoting reactions 
suggests some connection with the effect of water in promoting 
reactions between electrolytes, but the two cases seem to be very 
radically different. One chemist has gone so far as to say that 
probably no reaction can occur between two substances which 
are absolutely pure — but it is doubtful if such a generalization 
is justified. 

Carbon monoxide is very poisonous. Air containing one tenth 
of one per cent is distinctly dangerous if breathed for any length 
of time, and a smaller quantity, constantly present, would un- 
doubtedly cause chronic poisoning. It seems to combine with 
the hemoglobin of the blood and to alter it in such a way that 
the hemoglobin is no longer able to perform its proper function 
of combining with oxygen in the lungs and giving it up again for 
the oxidation of other substances in the tissues of the body. It 
acts as a cumulative poison, and recovery from its effects is often 
very slow. . 

Carbon monoxide may be considered as the anhydride of 
- formic acid, HCO.H (or H2CO:), which might also be called 
carbonous acid. At 200°-230°, best under pressure, it combines 
directly with sodium hydroxide to form sodium formate: 


CO + NaOH = HCO.Na 


The Cycle of Carbon in Nature. The carbon dioxide of the 
air furnishes the great storehouse from which carbon finds its 
way into living bodies through the growth of plants. These are 
able to use the energy of the sunlight for the reduction of carbon 
dioxide, possibly first to formaldehyde, CHO, which then unites 
with itself to form starch, (CgH105),. The nitrogen which is 
necessary to form proteins and other compounds and the other 
elements required for the growth of plants, especially compounds 
of phosphorus, potassium, sodium, calcium, magnesium, silicon 
and iron, must be furnished by the soil. The compounds which 


RESPIRATION CALORIMETER , 318 


are built up by the growth of plants are ultimately broken down 
and their carbon is returned to the air as carbon dioxide by one 
of three processes: 1. Wood or other vegetable material may be 
burned directly. 2. Vegetable substances may be used as food 
either by men or animals. After digestion and assimilation the 
carbon is sooner or later oxidized to carbon dioxide and is re- 
turned to the air, mostly through the lungs. 3. Vegetable or 
animal substances exposed to the action of bacteria, decay and 
the carbon is converted to carbon dioxide. 

Respiration Calorimeter. From the point of view of considera- 
tion of the transformations of energy, plants in their growth, 
by the reduction of carbon dioxide and the formation of com- 
pounds in which the carbon is combined with hydrogen and 
nitrogen as well as oxygen, store up energy received from the 
sun, in the form of combustible or edible carbon compounds. 
By burning and by the use of steam engines or other heat en- 
gines this stored chemical energy is transformed into heat energy 
or mechanical energy for practical use. When used as food by 
men or animals, the energy is also transformed both into heat 
and into muscular and mechanical energy. The amount of 
energy available in different kinds of food can be accurately 
determined by burning samples of them in a calorimeter (p. 25). 
In order to follow the transformations of energy which occur in 
the human body, the respiration calorimeter was devised by 
Professor W. O. Atwater with the aid of Dr. E. B. Rosa, and 
has been further developed by Benedict, Langworthy, and others. 
Figure 88 gives a cross section of one form of calorimeter which 
shows the general construction of the apparatus ; figure 89 shows 
an exterior view and some of the accessory apparatus for another 
form. The calorimeter consists of a small room in which a man 
may remain for experiments, which sometimes last for several 
days. The chamber is provided with tubes for the entrance 
and exit of the air necessary for respiration and also with pipes 
through which water can be circulated to take up the heat gen- 
erated by the body. The amount of food taken, any change in 
weight of the body which occurs, the amount of carbon dioxide 


314 A TEXTBOOK OF CHEMISTRY 




















DEAD AIR 
DEAD | AIR 




















—— 





































































































Seale: 1 Meter 


Fig. 88.— Horizontal cross-section of chair calorimeter, showing 
cross-section of copper wall at A, zine wall at B, hair-felt at H, 
and asbestos outer wall at F; also cross-section of all upright 
channels in the steel construction. At the right is the location 
of the ingoing and outgoing water and the thermometers. At 
C is shown the food aperture, and D is a gasket separating the 
two parts. The ingoing and outgoing air pipes are shown at 
the right inside the copper wall. The telephone is shown at the 
left, and in the center of the drawing is the chair with its foot- 
rest, G. In dotted line is shown the opening where the man 
enters. 


RESPIRATION CALORIMETER 315 


evolved and the quantity of heat given out by the body are all 
carefully measured. ‘The experiments involve, of course, many 
other details which cannot be given here. The investigations 
carried out with the calorimeter have demonstrated very 





Fig. 89 


clearly that the amount of heat generated when food is digested 
and oxidized in the human body is the same, within the limits of 
error of the experiments, as the amount of heat generated by 
burning the same quantity of food in a calorimeter. 

In some of the experiments the man in the calorimeter ex- 
pended muscular energy in driving machinery, using the pedals 
of a bicycle for the purpose. It was found that during this period 
the amount of heat given off by his body was largely increased. 
Thus, while at rest his body, in one experiment, gave out 112 
large calories per hour. When performing mechanical labor, 
on the other hand, 339 calories per hour were evolved, while 
the work done was 12,800 kilogrammeters, equivalent to 49 
calories, It will be seen from this that only about 18 per cent 


316 A TEXTBOOK OF CHEMISTRY 


of the additional expenditure of 276 calories required for the 
muscular work was actually converted into mechanical work. 
The human body, therefore, resembles a steam engine in dissi- 
pating as heat a large proportion of the energy required for its 
operation as a machine. 

It could not be shown by the respiration calorimeter that an 
appreciable amount of energy is required for mental work. 
In 22 experiments with persons performing mental work, as, for 
instance, in writing an examination or making arithmetical cal- 
culations, only one-half of one per cent more heat was given out 
from the body than that given out when the same individual 
was at rest. See Year Book of the United States Department 
of Agriculture for 1910, p. 307. 

The experiments seem to justify the conclusions that the 
compounds found in a living organism obey exactly the same 
chemical and physical laws as the same compounds outside of 
of the body. It cannot, however, be said that the difference 
between living and dead matter is simply a difference in physical 
and chemical properties. 

* Carbon Suboxide, C302, is formed when malonic acid, 
CH2(CO:H)s, is treated with phosphorus pentoxide, P2O;: 


OH HH OH 


‘ae ead Set 
O=C—C—C=0 —> O=C=C=C=0+ 2 HO 
Carbon Suboxide 


Carbon suboxide is a colorless liquid at a low temperature. It 
boils at 7° and has a very strong tendency to polymerize to a dark 
red substance, probably a mixture of compounds, having the 
same composition but evidently a much higher molecular weight. 

* Carbon Oxychloride, or Phosgene (Carbonyl Chloride) COC. . 
When a mixture of equal volumes of carbon monoxide and chlo- 
rine are exposed to the action of sunlight, the two gases unite 
to form carbonyl chloride. It is often called phosgene from this 
method of preparation (from ¢4s, light and yevvéw, to pro- 
duce). It is a colorless gas which may be easily condensed to 


CARBON BISULFIDE 317 


a liquid that boils at 8.2°. Carbonyl chloride is related to car- 
bonie acid just as sulfuryl chloride, SO2Cl., is to sulfuric acid. 


oie a OH oe 
oO) £70 
Ce) 2 ae Sen See ES ous Ive 
\e C] Se Se 
O—H OH Cl 
Carbonic Carbonyl Sulfuric Sulphuryl 
Acid Chloride Acid Chloride 


As the chloride of an acid, it is hydrolyzed by water: 
COCI, + 2 HO — 2 HCl + H:CO;— CO, + H:O 
With ammonia it gives the amide of carbonic acid, urea: 


Cl NH, 
CLO + 4NH, = c€o + 2NH,Cl 
Cl NH; 


Urea 


Carbon Bisulfide. When sulfur vapor is passed over heated 
charcoal, the two elements combine to form carbon bisulfide, 
CSs, which is now manufactured in electric furnaces. It is a 
volatile, inflammable liquid, which boils at 46.25°. It takes fire 
so easily that the vapors ignite when a glass rod, which has been 
warmed gently, is held over a dish containing the liquid. As a 
mixture of the vapor with air is explosive, this property makes 
extreme care necessary in factories where the substance is used. 
Carbon bisulfide is a good solvent for fats and india rubber, also 
for ordinary phosphorus, sulfur and iodine. It is used in vul- 
canizing india rubber, in the preparation of rubber cements and 
in extracting grease from wool. It is poisonous and is sometimes 
used to kill rats and ground squirrels. 

* Sulfocarbonates. Carbon bisulfide dissolves in solutions 
of alkaline sulfides, forming sulfocarbonates exactly as carbon 
dioxide forms carbonates with alkaline hydroxides : 

2 KOH + CO, = K2CO; + H:0 


Potassium 
Carbonate 


KS a Ss = K.CS3 


318 A TEXTBOOK OF CHEMISTRY 


or 2KHS+ CS. = K.CSs; + HS 
Potassium 
Sulfocarbonate 

Potassium sulfocarbonate is a yellow salt, easily soluble in 
water. It is used to destroy the phylloxera which sometimes 
cause great damage to grapevines. 

* Sulfocarbonic Acid, H2CSs3, separates as an oily liquid when 
hydrochloric acid is added to a solution of a sulfocarbonate. It 
soon decomposes into carbon disulfide and hydrogen sulfide. 


K.CS3 + 2 HCl = 2 KC] + H2Cs; => Hes + CS, 


The use of sulfocarbonates as germicides depends on a similar 
decomposition by means of the carbon dioxide of the air. 

* Carbon Oxysulfide, COS, is a compound intermediate be- 
tween carbon dioxide, COs, and: carbon bisulfide, CS2. It may 
be formed by the direct union of carbon monoxide, CO, and sul- 
fur, but it is best prepared by the decomposition of a thiocyanate 
by means of an acid: 


KCNS + HCl = KCI + HCNS 


Potassium Thiocyanic 
Thiocyanate Acid 
Hose Ce = NH; + 0O=C=5 


or KCNS + 2 HCl +H.O = COS + KCI + NH.Cl 


Carbon oxysulfide is a colorless, odorless gas which may be 
condensed to a liquid that boils at — 46.5°, a boiling point inter- 
mediate between those of carbon dioxide and carbon bisulfide, 
but much nearer to the boiling point of the former. It burns in 
air to carbon dioxide and sulfur dioxide. It is hydrolyzed by 
water to thiocarbonic acid, which then decomposes into carbon 
dioxide and hydrogen sulfide : 


Jou git 
O=C—p-5 HOH > O=CC —> O=C=0 + aX 
OH Na! 


Thiocarbonic 
cid 


OXIDES AND SULFIDES OF CARBON 319 


Cyanides. Potassium cyanide, KCN, may be formed by the 
direct union of the three elements at a high temperature, as when 
nitrogen is passed over a mixture of carbon and potassium car- 
bonate at a white heat : 


K,CO3 + 4 C + Ne = 2 KCN + 3 CO 


If organic matter containing nitrogen is heated with potassium 
carbonate and iron, a somewhat similar reaction occurs at a com- 
paratively low temperature, with the formation of a double 
cyanide of ferrous iron and potassium, called potassium fer- 
rocyanide, K,yFeCsNe (or Fe(CN)s.4 KCN). Cyanides are 
formed in a similar manner in the manufacture of illuminating 
gas and pass over into the ammoniacal gas liquors, from which 
they are recovered commercially. 

Hydrocyanic Acid or Prussic Acid, HCN, may be obtained by 
distilling a solution of potassium ferrocyanide or potassium cya- 
nide with dilute sulfuric acid. It is a liquid at low tempera- 
tures, which boils at 26.5°.. Hydrocyanic acid and its salts are 
among the most violent poisons known and should always be 
handled withextremecare. Theacid hasacharacteristic odor, but 
some persons seem to be unable to perceive it, and such individuals 
need to be especially careful in handling the acid or cyanides. 
Hydrocyanic acid is so weak an acid that it is liberated from its 
salts by the carbon dioxide of the air, and the simple cyanides 
have the odor characteristic of the acid. 

A dilute solution of hydrocyanic acid is sometimes used in 
medicine. 

Potassium Cyanide, KCN, or rather a mixture of potassium 
and sodium cyanides, is manufactured commercially by heating 
potassium ferrocyanide with metallic sodium : 


K,FeC,N, + 2 Na = 4 KCN + 2 NaCN + Fe 


The salt is used for many purposes in the laboratory, es- 
pecially in the preparation of organic compounds. It is also 
used for the extraction of gold from its ores and in silver plating. 

Complex Cyanides. The cyanides of the heavy metals, such 
as silver, iron, zinc, etc., are most of them insoluble in water, but 


320 A TEXTBOOK OF CHEMISTRY 


many of them will dissolve in a solution of potassium cyanide, 
ICCN. In the solutions obtained in this manner the atoms of 
silver or iron no longer conduct themselves as ions.! They will 
give no precipitate with reagents which precipitate them from 
their ordinary salts, as sodium chloride for silver or ammonium 
sulfide for iron. When an electric current passes through such 
, solution, the silver or iron travels with the cyanide ion toward 
the anode, while only the potassium goes toward the cathode. 
Finally, if the solutions are evaporated, definite crystalline 
compounds, potassium argenticyanide, KAgC2Ne, and potas- 
sium ferrocyanide, K4FeCsN,z.3 H2O, can be obtained. These 
facts indicate that such solutions contain complex argenticyanide 
(AgC.N._) and ferrocyanide (FeCgN¢g—) ions which hold together 
in solution as the elements of the sulfate (SO,7) or nitrate (NO3_° ) 
ions do. In further agreement with this interpretation, such 
solutions give with various solutions precipitates containing the 
characteristic complex group. <A considerable number of fer- 
rocyanides, for instance, may be obtained in this way. 

Ferric salts give with potassium cyanide, potassium ferricy- 
anide, K3FeCsN¢ (or FeC3N3.8 KCN), a salt which forms red 
crystals, while the crystals of the ferrocyanide are yellow. The 
ferrocyanide may be easily oxidized to the ferricyanide and the 
latter can be reduced to the ferrocyanide : 


2 K.FeCgNe ++ O + H,O =e K3FeCyNg + 2 KOH 
K3FeCeNeg -- H -- KOH = Ky4FeCeNe -} H.O 


Potassium ferrocyanide gives with ferric salts a blue precipi- 

tate of Prussian blue: 
3 KjFetG.Ne ob 4 Fe'MCls —- Fees (Fel! CeNe)3 aa 12 KCl 
Prussian Blue 

Potassium ferricyanide gives a similar deep blue precipitate 
with ferrous salts, but this is of variable composition and usually 
contains potassium. 

1A very small number of ions of these metals are doubtless 


present. In some solutions of this type hydrogen sulfide will give 
a precipitate, owing to the extreme insolubility of the sulfide formed. 


= 


§ 
3 
. 





COMPLEX CYANIDES Sok 


The ferric ferrocyanide is decomposed by potassium hydroxide 
with separation of ferric hydroxide while potassium ferrocyanide 
passes into solution. 

When a smaller amount of ferric chloride is added to a solution 
of potassium ferrocyanide a deep blue solution containing potas- 
sium ferric ferrocyanide is obtained : 

KiFe"CeNeg + Fe™Cls = KFe™Fe"C, Ng + 3 KCl 


Potassium Ferric 
Ferrocyanide 


In this solution a ferrous salt gives a precipitate of Turnbull’s 

blue: 

2 Khe Fe! CeNe -+- Eel Cl = Fe" (Fe™ Fel! CeNe)2 —+ 2 KCl 
Turnbull’s Blue 

The empirical formula of Prussian blue is Fe7CigNis, while 
that of Turnbull’s blue is FesCizgNi2. The empirical formula of 
ferrous ferricyanide is also FesCy2N12. 

Both Prussian blue and Turnbull’s blue are used as blue pig- 
ments and for “ blueing ” for laundry purposes. 

A solution of potassium argenticyanide, KAgC,Ne is used in 
manufacturing silver-plated ware. While the silver is trans- 
ferred toward the anode through the solution, it is also deposited 
on the cathode, and the fact that there are very few silver ions 
in the solution for some reason causes a smooth, coherent de- 
posit, while the silver deposited from a solution of silver nitrate 
usually assumes a crystalline form. From the silver anode, 
which is used in the electrolysis, silver passes into solution, re- 
generating the potassium argenticyanide which is decomposed 
at the cathode. 

* Potassium Cyanate, KCNO, may be prepared by heating 
potassium cyanide with lead oxide. It is very poisonous : 


KCN + PbO = KCNO + Pb 


* Potassium Thiocyanate, KCNS, is formed when a mixture 
of potassium cyanide and sulfur is heated. Potassium thio- 
cyanate and ammonium thiocyanate, NHiCNS, are used in 
testing solutions for the presence of iron in the ferric form because 


B22 A TEXTBOOK OF CHEMISTRY 


of the intense red color of solutions containing ferric thiocyanate, 
Fe(CNS)s. 

Cyanogen, C.N2. When mercuric cyanide, HgC2Ng, is heated, 
it decomposes into mercury and cyanogen, C2Ne, just as mercuric 
oxide decomposes into mercury and oxygen. Cyanogen is a 
colorless, poisonous gas, which burns with a characteristic pink 
flame. 


EXERCISES 


1. Carbon dioxide melts at — 56.4° under a pressure of 5.1 atmos- 
pheres. Under a pressure of 515 atmospheres (500 kg. per sq. em.) it 
melts at — 47.4°. Is the density of solid carbon dioxide greater or less 
than that of the liquid when both are present together? (See principle 
of van’t Hoff-Le Chatelier, p. 111.) 


2. Air contains, normally, about 3 parts in 10,000 of carbon dioxide. 
In accordance with Henry’s law what weight of the gas will be absorbed 
by one liter of water in contact with such air at 20°. (See p. 309 for the 
solubility of the gas in water at this temperature.) 


3. What volume of normal sodium hydroxide will be neutralized with 
formation of sodium bicarbonate, NaHCQOs, by one liter of water which 
has been in contact with ordinary air ? 


4. An effervescent drink is sometimes prepared by mixing two solu- 
tions containing sodium bicarbonate and cream of tartar (HKC,H,O,). 
What is the equation for the reaction ? 


5. A water contains 0.130 g. per liter of calcium sulfate, CaSQy. 
How may grams of crystallized sodium carbonate (NasCQ3.10 H2O) 
per cubic meter will be required to soften the water? How many grams 
of sodium fluoride ? 


6. How many grams of potassium cyanide will be required to reduce 
15 g. of stannic oxide, SnOs, to metallic tin ? 


CHAPTER XIX 


ALCOHOLS, ALDEHYDES, KETONES, ACIDS, FATS, CARBO- 
HYDRATES 


THE most important classes of compounds of carbon with 
oxygen and hydrogen are given in the heading of this chapter. 
The number of such compounds is very large and even a super- 
ficial knowledge of them can be gained only by a study of their 
structure, that is, by learning the arrangement of the atoms 
within their molecules. This is especially true because many 
different compounds having the same composition are known. 
Thus there are no less than seventy-five compounds having the 
formula C7;H;,02. Compounds like these which have the same 
composition but different properties are called isomers. The 
empirical formula of such a compound will give very little infor- 
mation about its properties, but a structural formula often shows 
to a chemist, at once, many important relationships. 

- Structural formulas are established mainly on the basis of 
three very simple principles : 

1. Valence. The valence of the elements. Carbon is almost 
always quadrivalent, oxygen is bivalent and hydrogen univalent. 
Nitrogen may be trivalent or quinquivalent. An example 
of the use of the principle of valence in explaining and predicting 
the structure of the hydrocarbons has been given (p. 284). 

2. Radicals. Groups of atoms called radicals hold together 
and retain their order of arrangement in passing from one com- 
pound to another. Thus in the reaction represented by the 


equation : C.H,O + HI = C.H;I + HOH 
Ethyl Ethyl 
Alcohol Todide 


the group, or radical, ethyl, C2Hs, is supposed to retain the same 


arrangement of its atoms in ethyl iodide that it has in ethyl 
323 


324 A TEXTBOOK OF CHEMISTRY 


alcohol and on the basis of this reaction we usually write the 
formula of ethyl alcohol as C2.H;OH. 
On the other hand methyl ether, which has the same composi- 
tion as ethyl alcohol, reacts with hydriodic acid thus : 
C.H,O + HI = CH;I + CH;0H 
Methyl Methyl Methyl 
Ether Iodide Alcohol 
On the basis of this reaction we give to methyl ether the 
formula CH;—O—CH;. Expanding these formulas, using the 
principle of valence, we have: 


Hea H H 
isagys | | 

ff yi By act and PO (a ami aaa 
| 
ja bodel H H 
Ethyl Alcohol Methyl Ether 


3. Substitution. When an atom or group leaves a compound 
and another atom or group enters it, the group which enters takes 
the place which was occupied by the atom or group which has left. 
In the reaction between methyl ether and hydriodic acid given 
above the group CH;0O leaves the compound and an iodine atom 
takes its place. 

Neither of the last two principles is universal in its application 
and there is some uncertainty as to the valence of the elements 
in some compounds, but, in spite of this, it has been possible to 
determine the structure of very many carbon compounds with 
practical certainty. It is impracticable within the limited space 
which seems suitable for this book to give the basis for the struc- 
ture assigned to the compounds mentioned, and in many cases 
only empirical formulas will be used. 

Alcohols. Alcohols may be defined as hydrocarbons in which 
one or more hydrogen atoms have been replaced by hydroxyl, OH. 

Methyl Alcohol, CH;OH, is obtained along with acetic acid 
and acetone by the destructive distillation of wood, and in an 
impure form it is often called wood spirit or wood alcohol. It is 
now used chiefly as an addition to ordinary alcohol to denaturize 


a 


oe 


ALCOHOLS B20 


it. (See below.) It is very much more poisonous than ordi- 
nary alcohol and fatalities have often occurred from drinking 
it or from breathing its vapor when varnishes made with it 
were used in a confined space. Blindness sometimes results 
from drinking or breathing it. Methyl alcohol boils at 66° and 
has a specific gravity of 0.7931 at 15.6°. 

Ethyl Alcohol, C.H;OH, is formed by the fermentation aaah 
yeast of liquids containing either ordinary cane sugar, as the 
juice of grapes or apples, or diluted sirups from the manufacture 
of sugar, or maltose or glucose, sugars formed by the action of 
malt on the starch of Indian corn or potatoes : 


C12H2O011 + H,O == 4. C.H;0H fe 4 CO, 
Maltose Alcohol 


Only dilute alcohol can be obtained by fermentation and strong 
alcohol is prepared by fractional distillation, the boiling point of 
alcohol being 78.3°. Absolute alcohol, or alcohol free from water, 
is prepared by allowing concentrated alcohol to stand with 
quicklime, CaO, which combines with the last portions of the 
water. The action of the lime may be hastened by warming 
the mixture. The specific gravity of absolute alcohol is 0.7933 
at 15°, referred to water at 4°. 

The most common beverages containing alcohol are beer 
(3.5-7.5 per cent of alcohol by volume), wine (9-17 per cent), 
cider (3.5-7.5 per cent), brandy, from the distillation of wine 
(44-55 per cent), whisky, from the distillation of alcoholic 
liquids made from grains (46-55 per cent) and rum (30-50 per 
cent). The per cents by weight are approximately four fifths of 
these. 

Alcohol is extensively used for burning, as a solvent, in mak- 
ing varnishes and in making pharmaceutical extracts, ether and 
other products. 

“ Denatured alcohol ” is an alcohol to which some substance 
has been added to render it unsuitable for use as a beverage. 
Wood alcohol, gasoline and bone oil are the most common 
additions. Such an alcohol is sold free of tax and may be used 


326 A TEXTBOOK OF CHEMISTRY 


for burning, for the manufacture of varnishes and for many other 
purposes. The substances added render it poisonous and unfit 
for drinking or for any medicinal use. 

Phenol or Carbolic Acid, CsH;OH, is obtained from coal tar 
and can be made artificially from benzene, CsHs. The pure 
compound is solid at ordinary temperatures but it liquefies on 
the addition of a small amount of water. It has been exten- 
sively used as a disinfectant and was the substance first used 
in antiseptic surgery. It is effective only when. applied in a 
solution. The boiling point is high and the vapor is never 
sufficiently concentrated to be of service. Pure phenol cor- 
rodes the tissues and is a violent poison. Phenol is the chief 
active constituent of the “ coal-tar dips’ 
sheep. 


2 


used in the care of 


CH,OH 
| 
Glycerol,! CHOH , or C3H;(OH)3, is an alcohol containing 


CH,0H 


three hydroxyl groups. It is obtained as a by-product in the 
manufacture of soap from fats (p. 332). When glycerol is 
treated with nitric acid under proper conditions, it reacts with it, 
forming glyceryl nitrate (nitroglycerin) : 


C3H;(OH); + 3 HNO; = C3H;(NOs3)3 aa 3 HOH 
Glyceryl 
Nitrate 
Nitroglycerin has a very much larger amount of chemical energy 
than the same elements combined in the form of carbon dioxide, 
water and free nitrogen, and under the influence of a denotating 
cap of fulminate of mercury the rearrangement of the atoms to 
the more stable, gaseous forms takes place so suddenly as to 
cause a violent explosion. It is used chiefly for blasting pur- 
poses. The liquid form is sometimes used directly, but usually 
1 Glycerol is commonly known as ‘‘glycerin.’’ The name glycerol 


is used by many writers and is to be preferred because the ending 
“ol” is used for names of alcohols in general. 


ALDEHYDES AND KETONES 327 


it is absorbed in sawdust or some other porous material and is 
then known as dynamite. 

The conduct of nitroglycerin illustrates in an unusually 
striking manner the fact that very many carbon compounds, 
perhaps even the majority of them, are in a condition of unstable 
equilibrium. The same elements may be combined in some 
other form which contains less chemical energy. This seems 
to be because the atoms do not separate and recombine so easily 
as is usual with the compounds of other elements. This slow- 
ness of reaction velocities and the fact that there are enormous 
differences in the velocity of different possible reactions, so that 
the compounds which are formed seldom represent the lowest 
content of chemical energy, are certainly very important factors 
in the formation of the extraordinary number of carbon com- 
pounds. 

Aldehydes and Ketones. By the oxidation of an alcohol 


z 


containing the group —C— , an aldehyde may be ob- 


Sita 
O 


tained. For this the characteristic group is —c€ , com- 
H 
bined with a single carbon atom. In a similar manner an 


alcohol containing the group Sa gives a ketone, which 
has the group »ye=0 combined with two carbon atoms. 


p? 
Formaldehyde, Me (or H:CO), is easily prepared 
H 


by the oxidation of methyl alcohol. A solution containing about 
40 per cent of formaldehyde is known, commercially, as for- 
malin and is much used as a disinfectant. For disinfecting rooms 
it must be disseminated in the form of a spray, or vaporized. 
It has the great advantage over sulfur dioxide, which was for- 
merly used for the same purpose, that it does not cause the dark- 
ening of metallic objects or injury to fabrics. It is a very power- 


328 A TEXTBOOK OF CHEMISTRY 


ful germicide and has been sometimes used as a preservative 
in milk and other articles of food. Such a use is very objec- 
tionable and is forbidden by law. Formaldehyde is a powerful 
poison and may produce painful wounds on the skin. 


® 
Benzaldehyde, CoHSCE , the chief constituent of oil of 
H 


bitter almonds and used as a flavoring extract, is one of the 
most common aldehydes. Another aldehyde is C%tral, 


CH, CHM 
'\C—CH—CH,—CH,—-C&CH~C& 
CH, H 


obtained from the oil of lemons and used as the starting point 
for the manufacture of lonone, an artificial substitute for the 
odor and the flavor of violets. 


BE 
LA is formed along with methyl alcohol 
CHs3 


and acetic acid by the destructive distillation of wood. It is 
also obtained by distilling calcium acetate or barium acetate. 
It is used in the manufacture of chloroform and as a solvent for 
acetylene. Acetone dissolves fats and is also used in the manu- 
facture of varnishes. 

Acids. The further oxidation of an aldehyde gives an acid, 


Ve 
for which the characteristic group is called carboxyl, Toe 


Acetone, 


The change from an alcohol to an acid is somewhat similar to 


the change from hydroxyl amine, NAH , to nitrous acid, 


\OH 


Y? 
N . In both cases the replacement of the two hydro- 
Sor 
gen atoms by an oxygen atom gives acid properties to the hydro- 
gen of the hydroxyl group. Very many acids are found among 
natural products and in articles of food. 


ACIDS 329 


| G0 
Formic Acid, H—C< , was first obtained from ants and 
O—H 


is, in part at least, the cause of the irritation from the sting of 
the bee. It is now manufactured on a large scale by absorbing 
carbon monoxide with sodium hydroxide at a temperature of 
190°-220° : 

NaOH + CO = HCOONa 


It is extensively used in making dyes and for other manu- 
facturing purposes as a substitute for acetic acid. 


YP 
Acetic Acid, CHs—CC (or HC.H;302), is the acid of 
OH 


vinegar. Itis formed by the action of Bacterium aceti or “mother 
of vinegar”’ on alcohol, either slowly, as with wine or cider 
stored in casks which are partly open to the air, or more rapidly 
by allowing an alcoholic liquid to trickle over wood shavings 
which have been inoculated with the bacteria. Good vinegar 
contains 4 per cent of acetic acid. Acetic acid is also formed 
by the destructive distillation of wood, and the crude acid ob- 
tained in this way is sometimes called pyroligneous acid. This 
crude acid, containing phenol and other substances which reén- 
force the antiseptic properties of the acetic acid, is sold 
under the name of “ liquid smoke ”’ as a preservative for meat. 


a 
escne NOE: 
Oxalic Acid, 0 (or H:C2Ox), is found in sheep sorrel and 
CF 
OH 


some other plants, but for technical purposes it is manufactured 
artifically by heating sodium formate with a catalyzer : 


CO.Na 

2HCONa = | +H, 
Sodium Formate CO.Na 
. Sodium 


Oxalate 


330 A TEXTBOOK OF CHEMISTRY 


Oxalic acid is a comparatively strong acid and is quite poisonous. 
It is used in calico printing and dyeing and in the bleaching of 
flax, straw and leather. 

The decomposition of oxalic acid into carbon dioxide, carbon 
monoxide and water when heated with concentrated sulfuric 
acid has been given (p. 811). The ammonium - salt, 
(NH4)eC204.H2O, is used to precipitate calcium because 
calcium oxalate, CaC2Ou, is only very slightly soluble. An 
acid potassium salt called potassium __tetraoxalate, 
KH3(C2Ox4)2.2 H2O, derived from the doubled molecule, is 
sometimes used as a standard for alkalimetry. 

Lactic Acid, CH;CHOHCO.H (or HC3H;O3), is formed by 
the fermentation of milk sugar and is found in sour milk. It is 
both an alcohol and an acid, as is apparent from its structural 
formula. 

CHOH—CO.H 
Tartaric Acid, | (or HsC4H4O.), is found in 
CHOH—CO,H 
grape juice. When the sugar of the juice ferments, forming 
alcohol, the acid potassium salt, KHC,H4sO., which is only 
slightly soluble in dilute alcohol, separates in a crude form called 
argol. From this the pure salt, known as cream of tartar, 
and the free acid are prepared. Cream of tartar is used in 
cooking and sometimes to aid in the formation of jelly from 
fruit juices. Mixed with sodium bicarbonate, NaHCQOs;, and 
flour or starch, it is used in making the more expensive kinds 
of baking powders. 

Potassium antimonyl tartrate, KSbOC4H,4Og, is called tartar 

emetic and is sometimes used as an emetic. 


CH,CO.H 
Citric Acid, C(OH)CO2H (or H3CgH;07), is found in the juice 
CH,CO.H 


of lemons and is the constituent which gives the sharp, sour 
taste to the fruit. 
A cold, concentrated solution of ammonium citrate, 


FATS BOL 


(NH4)3sCegH;O;, will dissolve some of the acid calcium phos- 
phates which are insoluble in water, while it will not dissolve 
tricalcium phosphate, Ca3(PO.)2. Its action in this respect 
is supposed to resemble the action of water containing carbon 
dioxide in the soil. Such a solution is used in the analysis of 
commercial fertilizers to distinguish between phosphates which 
are supposed to be readily available for the growth of crops and 
those which are only slowly available. 

Ammonium Ferric Citrate, (NH4),HFe(CsH;O7)o, is used in 
the preparation of “ blue-print ” paper. The paper is moistened 
with a solution containing a mixture of ammonium ferric citrate 
and potassium ferricyanide. When dry it is exposed to the 
action of light under a “ negative.’’ The two salts are easily 
soluble in water, but in the light the ferric salt or the ferricyanide 
is reduced by the citric acid. In either case, a blue, insoluble 
compound is formed (p. 320). On washing the paper with 
water the soluble salts are removed, while the portions exposed 
to the light remain a permanent blue. 

Benzoic Acid, CsH;CO:zH, is found in cranberries and in some 
other fruits. It is also manufactured by the oxidation of toluene, 
CsH;CHs3, from coal tar. Even in dilute solutions it destroys 
or prevents the growth of bacteria and it has been much used 
for this purpose. At the present time there is some difference 
of opinion among authorities as to whether small amounts 
of the acid taken with food have an injurious effect or not. Its 
use as a substitute for cleanliness or to disguise inferior products 
is, of course, condemned by every one.: 

Palmitic, Stearic and Oleic Acids, Fats. The natural fats, 
such as lard, tallow, butter, olive oil, cottonseed oil, linseed 
oil, and the oils found in nuts and cereals all contain compounds 
in which the hydrogen of three molecules of a monobasic or- 
ganic acid has been replaced by the trivalent radical glyceryl, 
C3H;. The most common of these fatty acids are palmitic 
acid, HCy6H3,Oo, stearic acid, HCjisH3;O2, and oleic acid, 
HC,sH3302. The corresponding compounds found in fats 
are palmitin, C3H;(CisH3102)3, stearin, CsHs(CisH35O2)3, and 


Oo2 A TEXTBOOK OF CHEMISTRY 


olein, C3H;(CisH3302)3. All three of these are found in lard. 
Stearin is found especially in tallow, palmitin in palm oil, olein 
in olive oil. The substance known commercially as stearin 
is in reality an impure stearic acid. It is used in laundries and 
for the manufacture of candles. Stearic and palmitic acids are 
solid at ordinary temperatures. Oleic acid is a liquid. 

Soaps. When fats are heated with a concentrated solution of 
sodium hydroxide, they are decomposed, forming a sodium salt 
of the fatty acid, and glycerol: 


C3Hs5(CigH3502)3 + 3 NaOHe,3 NaCisH3,02 ote C3H;(OH)3 
Stearin Sodium 
Stearate 


This process is called saponification and the sodium salts form 
the chief constituents of the ordinary soaps. The sodium salts 
formed by the saponification are separated from the aqueous solu- 
tion containing the glycerol by the addition of a concentrated 
solution of salt, in which the salts of the fatty acids are nearly 
insoluble. From the aqueous solution the glycerol is recovered 
by evaporation and distillation under diminished pressure. 

The action of soap depends on the fact that water containing 
soap in solution readily forms an emulsion with oily or greasy 
substances and so aids in their removal from fabrics or from the 
skin. 

The calcium salts of the fatty acids are mostly so insoluble 
in water that they cannot act in this way, and when soap is 
used with a water containing calcium salts in solution it can 
have little effect till enough of the soap has been used to pre- 
cipitate all of the calcium as calcium stearate, Ca(CigsH35O2)s, or 
in the form of similar compounds. ‘The separation of insoluble 
calcium salts when soap is used with hard waters is, of course, 
a familiar experience. 

Carbohydrates. A large class of organic substances found in 
plants is made up of compounds which contain oxygen and 
hydrogen in the same proportion as in water. On account of 
this composition these compounds are called carbohydrates, 


CARBOHYDRATES Boo 


but they do not contain oxygen and hydrogen in the form of 
water. Many of these compounds contain molecules with six, 
twelve, eighteen, or some multiple of six carbon atoms. The 
most important of the carbohydrates are the sugars, dextrins, 
starch and cellulose. 

Cane Sugar or Saccharose, Cy.H22011, is found in the juice 
of the sugar cane, in sugar beets, in the sap of the maple tree 
and in almost all sweet fruits and vegetables. From the sugar 
cané the juice is obtained by pressing the cane between heavy 
steel rollers. The. juice is concentrated under diminished 
pressure, to avoid the decomposition which would occur if the 
solution were boiled in the open air at atmospheric pressure. 
The partly concentrated solution is filtered through boneblack 
or animal charcoal to remove coloring matters, and is then evap- 
orated further till sufficiently concentrated so that the sugar 
will crystallize on cooling. The crystals are separated from 
the colored sirup by means of rapidly rotating centrifugal 
strainers. The manufacture of sugar from beets differs in 
many important details, but the general principles are the 
same. For maple sugar the juice is simply evaporated till 
the sirup is sufficiently concentrated to crystallize on cooling, 
the other substances present being of such a character as to 
give the sugar a desirable flavor. If purified by the methods 
described above, maple sugar would not differ from the sugar 
from sugar cane or sugar beets. 

Cane sugar crystallizes in well-formed monoclinic crystals, 
seen especially in rock candy. It melts at 160°, and if heated 
for a short time at that temperature or a little higher it is partly 
decomposed, giving a dark brown substance of indefinite com- 
position called caramel. 

A solution of cane sugar turns the plane of polarization of a 
ray of polarized light, which passes through it, to the right. 
The degree of rotation is almost exactly proportional to the 
concentration of the solution, and the measurement of the rota- 
‘tion in specially constructed polarimeters called saccharimeters 
is very much used as a basis for the control of operations in 


334 A TEXTBOOK OF CHEMISTRY 


sugar factories and for the collection of duty on sugar at ports 
of entry. | 

When warmed for a short time with a dilute acid, cane sugar 
is hydrolyzed to a mixture of glucose and fructose: 

CyeH22011 + H2O = CeH120¢6 + CoeH120¢ 
Glucose Fructose 

Glucose rotates the plane of polarized light to the right but 
fructose rotates the plane in a greater degree to the left, and the 
mixture has a levo-rotation. It is called for this reason invert 
sugar. The ease with which the hydrolysis takes place is a 
source of very considerable loss in the manufacture of sugar. 
A similar hydrolysis often occurs in fruit juices and in honey. 

Maltose, Ci2H».O1n, is formed along with maltodextrin by the 
action of the enzyme diastase (p. 344) on starch. Its formation 
is a very important step in the manufacture of alcohol from corn 
and other grains (p. 325). Maltose is hydrolyzed to glucose by 
the action of dilute acids. 

Lactose, or Milk Sugar, Cy.H22Q1;, is found in milk and can 
be obtained from whey as a by-product in the manufacture 
of cheese. It forms an important constituent of milk, as a food, 
and is used by preference, rather than cane sugar, as an addition 
to cows’ milk for feeding infants. 

Glucose, CsHi20¢, is formed together with fructose by the 
hydrolysis of cane sugar. It is also formed by the hydrolysis 
of starch with dilute acids: 

(CgH100s),, 7 H.O =a CeH120¢ 
Starch Glucose 

For the commercial manufacture dilute sulfuric acid is 
usually employed because on the subsequent addition of cal- 
cium carbonate the acid can be almost completely removed as 
the difficultly soluble calcium sulfate. Glucose, when pure, has 
about three fifths the sweetening power of the same weight of 
cane sugar. It is used in the manufacture of “ corn sirup,” in 
fruit preserves and in cheap grades of candy. The popular 
impression that glucose is harmful as an article of diet seems to 
have no experimental basis. Glucose is dextrorotatory and 


CARBOHYDRATES 335 


was formerly called dextrose, a name still used by some au- 
thors. 

In the disease called diabetes sugar and starch of the food 
which is eaten are changed to glucose and eliminated in the 
urine instead of being assimilated as they should be. The 
glucose can be detected by means of Fehling’s solution,! a solu- 
tion containing potassium sodium tartrate, KNaC4H.O,, copper 
sulfate, CuSO., and sodium hydroxide. Glucose reduces the 
copper of such a solution to cuprous oxide, Cu,O, which sepa- 
rates as a red precipitate when the mixture is boiled. 

Fructose, CsH120¢, is the second constituent of invert sugar 
(see above). It is levorotatory and was formerly often called 
levulose, but, since a second, exactly similar compound, which is 
dextrorotatory, is known, the designation fructose is preferred. 
Both glucose and fructose may be fermented to alcohol and 
carbon dioxide by the action of yeast. 

Starch, (CsHO;),, is found in the form of granules (Fig. 90) 
which differ very markedly in their organized structure but 
which, so far as is known, are identical in their chemical com- 
position. It is an important constituent of potatoes, Indian 
corn, rice, wheat, tapioca and many other cereals and vegetables. 
When the flour of a cereal is kneaded in a current of water, the 
fine starch granules float away, while most of the other con- 
stituents remain behind. ‘The residue consists chiefly of ni- 
trogenous substances and is called gluten. From the water carry- 
ing the starch granules in suspension the latter will settle out 
on standing or on allowing the water to flow slowly over tables 
very slightly inclined. The practical manufacture of starch in- 
volves, of course, many other details which need not be given here. 

Starch is the most important non-nitrogenous constituent of 
foods. The granules are covered with a thin coating which 


1 Fehling’s solution may be prepared by mixing equal volumes of 
two solutions containing: 1. 34.65 grams of copper sulfate 
(CuSO,4.5 H.O) in 500 ce. of water; 2. 173 grams of Rochelle 
salt (KNaC,H.0s.H.0) and 50 grams of sodium hydroxide in 
500 ec. of water. The copper of 1 ec. of the mixed solvtion will be 
precipitated by about 0.005 gram of glucose. 


336 A TEXTBOOK OF CHEMISTRY 


interferes with their digestion, and one of the most important 
effects produced by baking bread and cooking cereals and vege- 
tables is the bursting of the granules by the combined effect of 
heat and moisture. In the process of digestion starch seems to 





@ =U i) Dy (©@® 
B So «6 
02.90 0 p20 60D 19800029, D Ca2O2N.rr® 
Eosioro ncepseo® Warssahas ():080'0> 


3 oO v0, 0,° ° ° 0 
° oO ° oe » 
oo ie Oy 2, 4 Be» JC 3 = 
05.00 0 00. Oe dOQ\o e 4 
Bo BS 0 ¢ 
0.0 





° 
oO Oo ir woe Hey ° 
0° ve) oa Or ° °o ° ay a 
oo erro OV? ors) C) 


Fig. 90.—A, potato starch; B, rice starch; C, wheat starch 
(x 160). After Allen. 


be hydrolyzed to glucose, which is then used to form a part of 
larger molecules of compounds found in the tissues and fluids of 
the body. These compounds, in turn, are evidently easily avail- 
able in the animal economy for the production of heat and 
energy by their oxidation to carbon dioxide and water. 
Dextrin. If starch is moistened with very dilute nitric acid 
and heated for some time at 120°, it is converted into a soluble 
compound or mixture of compounds called dextrin. Other 
forms of dextrin may be obtained by the use of hydrochloric 
acid, by heat alone at a somewhat higher temperature, or by the 
action of the diastase of malt. The dextrins are more or less 
soluble in water and are used for the preparation of some kinds 


CARBOHYDRATES DOL 


of mucilage and for the backs of postage stamps and of gummed 
labels. 

Pectose, Pectin. Fruits of nearly all kinds, especially when 
not fully ripe, contain a substance called pectose, which is in- 
soluble in water but which when boiled with water in the 
presence of the fruit acids is decomposed or hydrolyzed with 
the formation of a soluble compound called pectin. Pectin 
forms a jelly with sugar in a slightly acid solution. The best 
conditions require enough sugar to give a solution boiling at 
about 103° and having a specific gravity while hot of 1.27 to 1.29. 
The fruit juice should contain not less than 0.5 to 0.7 per cent 
of an organic acid, calculated as tartaric acid. Usually an 
amount of sugar about one half to three fourths of the volume of 
the fruit juice is used. The boiling must not be long continued 
after separating the juice from the fruit, as this seems to destroy 
the pectin (see N. E. Goldthwaite, J. Ind. and Eng. Chem. /, 
333 and 2, 457; also Principles of Jelly Making, Univ. of IIl. 
Bulletin, Vol. 9, No. 36 (1912)). 

Pectin may be precipitated from fruit juices, which have 
been extracted by cooking, by means of alcohol, but its com- 
position and its properties as a definite compound have not been 
established. 

Cellulose. ‘The fiber of wood, flax, cotton, the outer coatings 
of cereals and many similar materials consist largely of an in- 
soluble substance having approximately the same composition 
as starch, (CsHiO;),. As coal and peat were formed from 
woody material, cellulose must be considered as the principal 
original constituent of all of our fuels except petroleum and 
natural gas. In hay and alfalfa it forms a very important con- 
stituent of the food of herbivorous animals. It also furnishes 
the basis for the manufacture of paper. The best grades of 
filter paper are nearly pure cellulose. 

Paper. The cheapest grades of paper are made from straw, 
the better grades from wood and the best from flax, or from 
cotton or linen rags. Many other fibrous materials may also 
be used. The materials are first treated with a variety of 


338 A TEXTBOOK OF CHEMISTRY 


chemicals to disintegrate them, bleach them and remove color- 
ing matters and other substances which are objectionable. 
This finally produces a thin, uniform pulp which can be spread 
out evenly to form the sheet of paper. In the best’ kinds 
of paper the fibers must remain as long and as strong as 
possible. It is also necessary to remove substances which 
turn brown on exposure to the light and which cause the 
paper to darken with age. 

The glazed surface of paper, necessary to prevent the absorp- 
tion and spreading of ink, is obtained by the application of rosin, 
aluminium sulfate and other substances as sizes. 

Gun Cotton, Celluloid, Lacquers. When cotton is digested 
with a mixture of concentrated nitric and sulfuric acids a variety 
of compounds called nitro-celluloses are formed, differing with 
the concentration of the acids used, the duration of the treatment 
and the physical condition of the material. The most highly 
nitrated product has the composition Cy.H140.(NOs3)., and 
is called hexanitrocellulose. These products are powerful ex- 
plosives and are used in torpedoes and also as the principal 
constituent of smokeless powder. 

Less highly nitrated forms dissolve in a mixture of alcohol 
and ether, forming collodion. The evaporation of the solvent 
leaves the material as a thin, elastic film which is used to hold 
the silver compound for the wet plates in photography. It is 
also sometimes used to protect wounds from the access of 
bacteria. Similar products dissolved in amyl acetate form 
excellent lacquers. Mixed with or dissolved in camphor they 
form celluloid. 


CHAPTER XX 


AMINES, DYES, ALKALOIDS, PROTEINS, ENZYMES, FOODS 
AND NUTRITION 


IF one or more of the hydrogen atoms of ammonia are replaced 
by organic radicals, a great variety of compounds called amines 
are formed. These compounds combine directly with acids to 
form salts and are often called organic bases, but it should be 
remembered that the true bases are related to the amines in 
the same way that ammonium hydroxide, NH,OH, is related 
to ammonia, NH;. The “strength” of these bases varies 
greatly according to the nature of the radicals which they con- 
tain. Thus methyl amine, CH3;NHbp, forms in aqueous solution 
a much stronger base (CH;NH3;0H) than ammonium hydroxide, 
while aniline, CsH;N Hp, gives a very much weaker base. Both, 
however, form well-crystallized, definite salts, as methyl am- 
monium chloride, CH3NH3Cl, or aniline hydrochloride, 
CesHsNH3Cl. 

Methyl Amine, CH;NH2. Ammonia combines directly with 
methyl iodide, CH3I, to form methyl ammonium iodide, 


NES H 
CH3I + NH; = H SNC 


Methyl Ammonium Iodide 


When methyl ammonium iodide is warmed with a concentrated 
solution of sodium hydroxide, methyl amine escapes as a gas, 
exactly as ammonia escapes when ammonium chloride is treated 
in the same way: 


339 


340 A TEXTBOOK OF CHEMISTRY 


Methyl amine is a gas with a disagreeable odor resembling 
that of herring brine. It resembles ammonia in its general 
properties, but is more easily combustible. 

Aniline. By treating benzene, CsH¢, with nitric acid a com- 
pound called nitrobenzene, CsH;NOs, can be prepared. When 
this compound is reduced by tin and hydrochloric acid, or, 
commercially, by iron and acetic acid, aniline, CsH;N Hg, is ob- 
tained. When pure, aniline is a colorless oil which boils at 184°. 
It is made on a large scale for use in the manufacture of a great 
variety of dyes and for the preparation of several valuable com- 
pounds used in medicine, especially of acetanilide (antifebrin), 


CH 





CO 
CHNHCsHeO mann aeatin crite CoHs—NC 
N(CH3)—C—CHs. 
NHC,H30 
Phenacetine, CHC , may also be considered as a 
OC2H; 
derivative of aniline. 

Dyes. ‘Till the middle of the nineteenth century all of the 
dyes used for coloring fabrics were either inorganic compounds 
or were natural, vegetable or animal products. Vegetable 
products were chiefly used, but the number of those available 
was comparatively small, the two of greatest importance being, 
probably, indigo and alizarin, or Turkey red. In 1856 Sir 
William Perkin, in the course of some experiments which he 
tried in the hope of obtaining quinine from aniline, discovered a 
beautiful violet compound, mauve, which can be manufactured 
by the oxidation of aniline. During the next few years he estab- 
lished the manufacture of mauve on a commercial basis. This 
proved to be the starting point for a great industry for the prep- 
aration of thousands of different colors by synthetic processes. 
Some of the artificial dyes are identical with those obtained 
from vegetable sources. Many others rival these in brilliancy, 
in being insoluble or “ fast ”’ when the fabrics dyed with them 
are washed and in resisting the action of light. Others are not 
so good as some of the natural dyes in these last particulars. 
Colors of almost every conceivable shade are now available. 


DYES d41 


Alizarin, CyHs0.(OH)». Shortly after the discovery of 
mauve, Graebe and Liebermann, two German chemists, showed 
that by distilling alizarin, the coloring matter of Turkey red, 
with zine dust, it is reduced to anthracene, a hydrocarbon 
found in coal tar: 


CyH,O.(OH)> 4e a Zn a H,O = Cy4H i +- 5 ZnO 


Alizarin Anthracene 


Soon after this, methods were developed for the manufacture 
of alizarin from anthracene, and in a very few years the artificial 
product displaced the natural dye and the raising of madder 
root, from which the dye had been obtained, practically ceased. 

Indigo, CigHiO2.N2. This dye, which has been extensively 
used for many centuries, has been obtained until recently almost 
exclusively from a plant growing in India. In 1881 Professor 
Baeyer in Munich discovered a method of making indigo arti- 
ficially. The process was complicated, however, and the original 
material used, toluene, was too expensive to allow of the profit- 
able manufacture of the dye. Twenty years of continuous 
study in University laboratories and in factories were required 
before a successful process was developed. One process used 
starts with naphthalene, CioHs, and acetic acid as the original 
materials. In 1901 it was so far developed that the Badische 
Soda-Anilin Fabrik had been willing to spend $4,500,000 in pre- 
paring for the manufacture on a large scale. Since then the 
amount of the synthetic indigo produced by this and other 
processes has increased each year, and it has, now, largely 
displaced the natural product. 

Indigo is extremely insoluble in almost all solvents which do 
not change it chemically, and its value depends very largely on 
this property, which makes it a “fast” color. In order to fix 
it on the fiber it is dissolved by reducing it with ferrous hydroxide 
in the presence of calcium hydroxide or, of recent years, with 
an alkaline solution of sodium hyposulfite, Na2S,0,. The in- 
digo white formed by the reduction is a weak acid and forms a 
soluble salt with the calcium or sodium. Fabrics which are to 


342 A TEXTBOOK OF CHEMISTRY 


be dyed are dipped in the alkaline solution. On exposure to the 
air the indigo white is oxidized back to indigo, which remains 
firmly attached to the fiber : 


CigHpO2Ne 4- 2 Fe(OH). a 2 H.O 
Indigo 
= H.CigH O2N2 -+- 2 Fe(OH); 
Indigo White 


H.C1gH19O2N2 fe Ca(OH). — CaC gH pOoNe oa yi H,O 
Soluble Calcium Salt 
of Indigo White 


Mordants. Some dyes, as indigo, are so insoluble that if 
once formed in contact with the fiber of a fabric, they will not 
dissolve and they produce a “ fast” color. Other dyes seem to 
combine with fibers directly to form insoluble compounds. 
Such dyes are called “ substantive’ dyes because they are in- 
dependent of the use of other substance required to fix them 
on the fiber. Substantive dyes which may be used for silk or 
wool are much more common than those for cotton or linen. 
Other dyes, which are called “ adjective’ require a mordant, 
with which they form an insoluble compound, to fix them. 
The most common mordants are aluminium acetate, ferric 
acetate, potassium dichromate and tannic acid. 

Alkaloids. ‘There is a considerable number of nitrogenous 
compounds, found in plants, which combine with acids to form 
crystalline salts in the same way that the amines do. Some of 
these are comparatively simple amines, but most of them are 
complex in their structure. Many of them have some very 
marked physiological action as poisons or as medicines. Many 
which are poisonous are used as medicines in small doses. 

Nicotine, C1H,4No, is a colorless oil found in tobacco, which 
contains from 2 to 8 per cent of the alkaloid. It is very poison- 
ous. 

Coniine, CsH;;N, the alkaloid of hemlock, is also a liquid. 
It is historically interesting as the active principle of the fatal 
draught taken by Socrates. 


ALKALOIDS. PROTEINS 343 


Atropine, C17H2303N, is found in Atropa belladonna. It is 
used to dilate the pupil of the eye and is an active poison. 

Cocaine, C,;H»,0,N, is found in coca leaves. It is used to 
produce local anesthesia. A careful study of cocaine has 
shown that it is a derivative of benzoic acid and the group 
derived from that acid is chiefly effective in giving to it its 
valuable qualities. On the basis of this discovery other alka- 
loids having, in part, a similar structure have been prepared. 
Some of these retain the anesthetic effect of cocaine and are 
less poisonous. 

Morphine, C,7H1903N.H2O, is the most important alkaloid 
of opium and is the chief constituent which gives to laudanum 
and paregoric their poisonous and sedative qualities. Paregoric 
also contains camphor and aromatic oils which may have as 
much effect as the morphine. Opium is obtained from the 
poppy. 

Quinine, CopH2sO2No, is obtained from Peruvian bark. It 
is a specific in malarial fevers. 

Strychnine, Co:H22O2Na, is found in Strychnos nux vomica. It 
is a violent poison, producing convulsions. <A dose of 0.06 gram 
is considered fatal. In small doses it is a powerful stimulant. 

Ptomaines. In the putrefaction of fish or meat under the 
influence of bacteria and sometimes in the putrefaction of vege- 
table substances, a variety of basic compounds called ptomaines 
is formed. Some of these are poisonous, and illness from 
ptomaine poisoning often results from eating spoiled meats, 
especially fish. A few of them give color reactions similar to 
those given by the vegetable alkaloids. Their presence often 
greatly increases the difficulty of identifying alkaloids in toxical 
analysis. 

Proteins. The most important compounds in both animal 
and vegetable organisms seem to be the proteins. These con- 
tain carbon, hydrogen, oxygen and nitrogen, usually sulfur and 
sometimes phosphorus or iron. The albumen of the white of an 
egg is a nearly pure, typical protein. Proteins form the larger 
part of muscular fiber, of the casein of milk and of the gluten of 


344 A TEXTBOOK OF CHEMISTRY 


flour. The molecular weight of the proteins contained in the 
substances just mentioned is very high — as much as 15,000, at 
least. When foods containing proteins are eaten, during the 
process of digestion they are hydrolyzed, under the influence of 
the enzymes (see below) of the digestive fluids, giving less com- 
plex proteins called albumoses, and by further hydrolysis 
amino acids. These pass into solution and so into the circula- 
tion of the blood and from them and from other portions of the 
food the organism reconstructs the proteins which enter into 
‘the tissues and fluids of the body, replacing those proteins 
which are constantly being oxidized to furnish heat and energy 
for the organism. In part, they are oxidized directly without 
being transformed into tissues. 

Enzymes. The larger portion of the foods which are eaten 
are insoluble in water and in a form which could not be directly 
assimilated. In the course of the digestive tract is found a 
series of substances, called enzymes, which act upon the food 
catalytically, hydrolyzing it or changing it so that it becomes 
soluble, and emulsifying the fats. The most important of these 
enzymes are the ptyalin of the saliva, which changes starch to 
sugar, pepsin of the gastric juice of the stomach, which, with the 
aid of hydrochloric acid, normally present, changes the proteins 
to albumoses and renders them soluble, and trypsin from the 
pancreas. The fluids of the digestive tract are alternately 
alkaline and acid. 

Many other enzymes are known. One of the first to be dis- 
covered was the dzastase, which is formed during the germina- 
tion of barley and which forms the active constituent of malt. 
It transforms the starch of grains into maltose and dextrin in 
the manufacture of alcohol. Yeast cells contain an enzyme, 
zymase, which transforms glucose, fructose or maltose to alcohol. 

Toxins, Antitoxins. In snake venom and in a few plants, 
especially in the castor bean, substances are found which seem 
to resemble the enzymes, but which produce disastrous, poison- 
ous effects upon the animal organism. These are called toxins. 
Toxins seem also to be formed by the action of bacteria in cer- 


FOODS AND NUTRITION 045 


tain diseases and doubtless their formation is often a chief factor 
in the progress of the disease. It has been discovered that 
animals subjected to the effect of such a toxin develop a sub- 
stance which appears to combine with it and render it harmless. 
By use of this principle it has been possible to prepare antitoxins 
which are powerful agents in combating these diseases. 

Urea, CON2H,. About 85 per cent of the nitrogen taken as 
food is eliminated from the human body in the form of urea. 
This may be considered as formed by the union of carbon 
dioxide and ammonia, followed by the loss of water : 


NH, 
O OH NHe 
c€ + 2NH; = = c€o + 1,0 
O . OH NH, 


NHe 


Urea is also formed by rearrangement when a solution of 
ammonium cyanate is evaporated : . 


ye 

NH,—O—C=N —> Ce 
Ammonium Cyanate NH, 
Urea 


This transformation of ammonium cyanate into urea, which 
was discovered by Wohler in 1828, was the first synthesis of an 
“ organic”? compound from inorganic materials and it was the 
beginning of a long series of syntheses which have demonstrated 
that many compounds prepared in the laboratory are identical 
with the same compounds found in animals and vegetables. 

Nutrition. An adult man weighing about 70 kilograms, when 
on an average, mixed diet, eliminates from his system 16 to 18 
grams of nitrogen per day. This is equivalent to the consump- 
tion of 100 to 112 grams of digestible protein, which could be 
obtained from 3 to 3% liters of milk, 1150 to 1250 grams of 
white bread, 600 to 750 grams of fresh fish, or 500 to 560 grams 
of lean beef. From the experiments with the respiration calorim- 
eter (p. 313) it seems that in a room at 20° the body of an 


346 A TEXTBOOK OF CHEMISTRY 


adult weighing 70 kilograms loses about 2200 calories in 
24 hours. The protein referred to above would give by its 
oxidation in the body, only 420 to 475 calories.!_ To furnish the 
balance of energy required by the body, about 325 grams of 
carbohydrates (starch, sugar, etc.) and 50 grams of fat would be 
required. Theoretically it does not matter whether the extra 
energy 1s supplied by carbohydrates or by fat. Practically, 
both are usually taken somewhat in the proportions given. 

When engaged in muscular labor for 8 hours a day, about 20 
grams of nitrogen are eliminated and about 1800 addition calories 
are given out as heat and mechanical energy. ‘This would re- 
quire a total,of about 125 grams of protein, 625 grams of carbo- 
hydrates and 100 grams of fat. In the respiration calorimeter, 
where the man drove a stationary bicycle arranged to measure 
the work performed, about 20 per cent of the energy of the 
extra food required was converted into mechanical energy. 
The remainder was dissipated as heat given out from his body. 
This indicates that the human body, considered as a machine, 
is somewhat more efficient than the best steam engines. But 
the food required as the source of energy is, of course, much 
more expensive than coal. 

A large number of dietary studies have given results which 
indicate that the average American diet for an adult man 
weighing 70 kilograms is approximately that given in the table 
on opposite page. 

The consideration of the amount of protein and of heat energy 
required by the body, while undoubtedly of great importance 
in deciding upon the character of the diet which should be 
selected, takes account of only a few of the factors which ought 
to be considered. About many of these factors our knowledge 
is still very imperfect. It is claimed by some writers that the 
human body may be maintained in a state of health with the 

1In the bomb calorimeter 1 gram of protein gives about 6.55 
large calories. When taken as a food a part of the protein is elim- 
inated in the form of uric acid, creatinine and other compounds, 


which may still be burned with evolution of heat, hence the heat of 
oxidation of protein in the body is only about 4.2 calories per gram. 


FOODS AND NUTRITION 347 


PRoporTIONS oF NuTRIENTS FURNISHED BY DIFFERENT Foop 
MATERIALS IN THE AVERAGE AMERICAN DIETARY 


























PRo- oe CaRBo- 
TEIN HYDRATES 
Animal foods : iE Ys - ( 
FLOral Ticats ime Ue wetee sy ane eae LO.0 29.7| 58.8 — 
Rist hes re ee LS oo 0 — 
ip asuart whe Leek elke A mes Meee ed 4 Te 2.9 — 
iDairye products), cyets. =) “a 18:4 10.0| 25.7 3.6 
Unclassified animal foods . .| 0.2 38.5 O26 a2 0.3 
Vegetable foods : 
otabecrealsee 0 mayen) 61 OU-0 43.0| 9.1) 61.8 
Sugar, molasses,etc. . . .| 5.4 — —| 17.6 
Legumes, tubers, vegetables . | 20.3 8.7 1.0 12.0 
Fruits, including nuts .°. .| 4.4 OMe SUE 4.3 
Unclassified vegetable foods .| 0.5 61.2 = — — 
Miscellaneous food materials . 0.3 Oe e)G 0.4 
100.0 | 100.0 | 100.0 | 100.0 























use of a very much smaller amount of proteins, indeed with 
about 60 per cent of that given above, and experiments have 
been tried which tend to support this point of view. It has 
been shown, too, that the proteins from different sources differ 
very much in the character of the amino acids formed by their 
hydrolysis. Inasmuch as certain amino acids in definite quan- 
tities are required to restore the wasted tissues of the body, it 
seems certain that some of the proteins are much more suitable 
than others for use as food. But inquiry along these lines is 
recent and has not proceeded far enough to lead to final con- 
clusions. 

Finally, there are many of the inorganic elements which are 
absolutely essential to the health of the organism, such as 
sodium, chlorine, iron, calcium, silicon, sulfur, phosphorus and 
even lodine. 


CHAPTER XXI 


SILICON, BORON, GERMANIUM, TIN, LEAD, TITANIUM, 
ZIRCONIUM, CERIUM, THORIUM 


THE atomic weights of the nonmetallic elements, in round 
numbers, arranged in the order of the groups of the Periodic 
System, are as follows. The most closely related metallic ele- 
ments are also given. All of the nonmetallic elements are above 
and to the right of the dotted line. 


B hee: 12) Neg te DL Gleae Feat 1D Ne. 20 


Ale 2 (aisiam2on 31 § YS ULM Sess at 40 


Ga 70 Ge 721 As, W75t Se P70 tBre (80k es 


In’ -114 Sn’ 1182:Sbe1209 Te 127-5 5 lanl 27 ce 


Tl 204 Pb 207 Bi 208 aoe! a Rn 222 


Silicon. Si, 28.06. Occurrence. Silicon is the most widely 
distributed and abundant element after oxygen. It forms 
about one fourth of that portion of the earth which we can 
examine. It is the characteristic element of minerals almost 
as much as carbon is the characteristic element of living matter 
— though there are many minerals which do not contain sili- 
con, and the number of silicon compounds is very much smaller 
than that of the carbon compounds. 

Silicon is found in nature in the form of the dioxide, SiOz, in 
rock crystal, the purest form of quartz, amethyst, jasper and 
agate, flint, or chalcedony. Silicon dioxide also forms the prin- 
cipal constituent of the sandstones and of ordinary sand. A 
hydrated dioxide, containing varying amounts of water, is called 
opal. Silicon is found in a great variety of silicates, all of which 
may be considered as salts of silicic acids, of which silicon dioxide 
is the common anhydride. Among these may be mentioned 
orthoclase, one of the feldspars, KAISi303, mica, KH2A13(SiO«)3, 

348 


SILICON CARBIDE 349 


kaolin, H2Al2(SiO4)2.H2O, a chief constituent of clay, asbestos, 
a variety of amphibole, CasMg3(SiO,)s, tale or soapstone 
(alberene), H»MgsSisOx, serpentine (meerschaum), HyMg3Si.Ox, 
garnet, Ca3Fe2(SiO,)s3, or Ca3Al2(SiOx)s3, topaz, Aly2SigQo5F 10, 
tourmaline, Al,BsO15.4 H2NaAl3(SiO4)3, and beryl, BezAle(SiO3)¢. 
Only a few ores of common metals are silicates, the most impor- 
tant being calamine, Zn2SiO,. 

Preparation. Silicon is never found free in nature. It is 
most easily prepared by heating a mixture of fine sand with 
magnesium : 


SiO. + 2 Mg = 2 MgO + Si 


The silicon obtained in this way is an amorphous brown powder 
insoluble in water and acids, except in a mixture of hydrofluoric 
and nitric acids. It dissolves in a solution of sodium hydroxide 
with evolution of hydrogen: 


S1 + 2 NaOH + H.O = Na,SiO; + 4H 


Sodium Silicate 


Silicon may be crystallized from its solution in melted zinc 
and then forms brilliant needles having a metallic luster. It is 
now made in electric furnaces for use in the steel industry. 

Hydrogen Silicide, SiH,. If fine sand is heated with four atoms 
of magnesium for each molecule of the silicon dioxide, the silicon 
combines with the magnesium to form magnesium silicide, 
Mg.Si. When this is treated with a dilute acid, hydrogen sili- 
cide is formed : 


Mgssi + 4HCl=2MeCl, + SiH, 
Magnesium Hydrogen 
Silicide Silicide 
Hydrogen silicide is a colorless gas which takes fire spontane- 
ously on coming to the air and burns to water and silicon dioxide. 
It may be condensed to a liquid, which boils at a very low tem- 
perature. 
Silicon Carbide. Carborundum, SiC. By heating silicon 
dioxide with coke in an electric furnace it may be reduced and 


300 A TEXTBOOK OF CHEMISTRY 


the silicon and carbon unite to form silicon carbide or carborun- 
dum, which, when pure, crystallizes in beautiful, colorless 
needles : 


SiO. + 3 C = SiC + 2 CO 


Carborundum is the hardest substance known except the car- 
bide of boron and the diamond, and it is manufactured for 
use as an abrasive. For this purpose it has partly displaced 
corundum, Al,O3, which is used under the name of emery. 
Carborundum is also used as a refractory material in the con- 
struction of furnaces and to remove gases from steel. 

Silicon Fluoride, Sify. The formation of silicon fluoride in the 
etching of glass by hydrofluoric acid has been referred to. The 
compound is most easily prepared by warming a mixture of sand, 
SiOz, calcium fluoride, CaF2, and concentrated sulfuric acid : 


2 CaF, + SiO. + 2 H.SO, =e CaSO, + Sif’, a x, H.O 


Silicon fluoride is a gas, but may be condensed to a solid which 
melts at — 97° and has a vapor pressure of 760 mm. at — 90°. 

Fluosilicic Acid, H.Sil's. Silicon fluoride is hydrolyzed by 
water, as would be expected of a halogen compound of a non- 
metallic element, but the hydrofluoric acid formed combines 
with some of the undecomposed silicon fluoride to form a com- 
plex acid, fluosilicie acid : 


Sif, + 4 HOH = Si(OH),+ 4 HF 
Silicon Silicie 
Fluoride Acid 
2HF+ Sif, = Hasiks 


Fluosilicic 
Acid 


The potassium salt of fluosilicie acid, K2Sik’s, is difficultly — 
soluble. The barium salt, BaSil’s, is also extremely insoluble, 
even in dilute acids. It is the only salt of barium likely to be 
mistaken for barium sulfate, BaSO., when barium chloride, 
BaCly, is used to test for sulfates in a dilute acid solution. 


SILICON DIOXIDE ta | 


Silicon Tetrachloride, SiCl4, is formed when chlorine is passed 
over heated silicon. It was formerly prepared by heating a mix- 
ture of silicon dioxide and charcoal in a current of chlorine: 


p102 + 2 Cl. + 2C = SiCl, + 2 CO 


This method of preparation, which was formerly much used 
to obtain chlorides of elements, such as silicon, aluminium and 
chromium, which cannot be reduced from their oxides by carbon 
at any moderate temperature, has ceased to be of practical 
importance since other methods have been developed for the 
preparation of the free elements. The process depends on the 
simultaneous use of chlorine and carbon to cause the separation 
of the silicon and oxygen. Silicon tetrachloride is a liquid 
which boils at 56.9°. 

It is hydrolyzed by water to silicic acid, Si(OH)., and hydro- 
chloric acid. 

Silicon Hexaiodide, Siels, is formed when silicon tetraiodide, 
Sila, is heated to 290°-800° with powdered silver. It crystallizes 
from carbon bisulfide in colorless prisms. It is hydrolyzed by 
water to silicooxalic acid : 

Sielg + 4 HOH = H.Si.0, + 6 HI 
Silicooxalic 
Acid 

These compounds and some others, which have been prepared, 
show that silicon atoms may combine together as carbon atoms 
do, but only a few such compounds are known and most of these 
are comparatively unstable. Silicooxalic acid, H2SiO., 1s so 
unstable as to be explosive. It seems evident that stable com- 
plex molecules containing silicon are formed only when the atoms 
are held together through the agency of some other element, 
such as oxygen. (See below under Silicic Acids.) 

Silicon Dioxide or Silica, SiO.. In addition to the forms of 
occurrence already given, infusorial earth or “ Kieselguhr,” 
a porous material composed of the skeletons of infusoria, may be 
mentioned. It is used to absorb nitroglycerin in the manu- 
facture of dynamite and as a packing material for bottles con- 


352 A TEXTBOOK OF CHEMISTRY 


taining bromine or other corrosive chemicals. It is also used 
in sapolio and in other scouring soaps and scouring materials. 

Clear specimens of rock crystal are sometimes used in the 
preparation of prisms and lenses which are more transparent than 
ordinary glass to ultra-violet light. Flint and pure sands and 
sandstones are used in the manufacture of glass. 

Silicon dioxide is found in nature in two forms. Rock crystal 
or quartz, which is much the more common, crystallizes in forms 
of the hexagonal system and has a specific gravity of 2.6. Quartz 
can be formed only at temperatures below 900°. Above that 
temperature it changes to tridymite, which crystallizes in the 
rhombic system and has a specific gravity of 2.28. 

Quartz melts at about 1750°, but softens and becomes plastic 
at 1600° or below. By means of the electric furnace or the 
oxyhydrogen flame it can be melted and fashioned into tubes, 
crucibles, beakers, flasks, thermometers and other laboratory 
utensils. It has the advantage over glass that it can be heated 
to very high temperatures without melting and also that it has 
such a small coefficient of expansion for a change of temperature 
that it does not crack when subjected to sudden heating or 
cooling. Partly for the same reason thermometers made from 
it show no depression of the zero point after use at high tempera- 
tures, as is common with glass thermometers. 

Fused quartz has a specific gravity of only 2.20, nearly the 
same as that of tridymite. Fused tridymite would doubtless 
be a more correct name than fused quartz. 

Artificial Silicates. When silicon dioxide is heated with so- 
dium carbonate or with the oxide or carbonate of almost any 
metal, a silicate is formed : 


NasCO3 + SiO, — Na.SiO3 + COE 
Sodium Sodium 
Carbonate Silicate 


CaO + SiO, = CaSiOs; 


The industrial importance of reactions of this character will 
be seen when it is stated that similar reactions are used for the 


SILICIC ACIDS BOD 


manufacture of glass, for the formation of fusible slags in the 
manufacture of iron and in other metallurgical operations and 
in the manufacture of Portland cement. 

The silicates of sodium and potassium are soluble in water, 
and the former, especially, is called water glass and is used as an 
addition to laundry soaps, for preserving eggs and for the fire- 
proofing of wood and fabrics. Almost all other silicates and 
mixed silicates are insoluble or nearly insoluble in water. 

Silicic Acids. If concentrated hydrochloric acid is added 
quickly to a solution of sodium silicate, the silicate seems to be 
completely decomposed, but the 
silicic acid formed does not pre- 
cipitate. If such an acidified 





solution is dialyzed by placing 
it in a parchment sack suspended 
in water (Fig. 91), the sodium 
chloride will diffuse through the 
walls of the sack, and by repeat- 
edly changing the water on the 
outside the chloride can be al- 
most completely removed, leaving Fie. 91 
a colloidal solution of silicic acid. 
Colloidal solutions may also be prepared by the hydrolysis 
of esters of silicic acid such as the methyl] ester, (CH3)4SiOx. 

According to the method of preparation silicic acid may be 
either a negative or a positive colloid. When it is a negative 
colloid it retains some anion, such as the chloride ion, Cl~, which 
cannot be removed by dialysis or washing. When it is a positive 
colloid it retains some cation as the sodium ion, Na‘, in the same 
way. Solutions of the first class are precipitated by solutions 
containing a bivalent cation, such as barium chloride, BaClk, 
while those of the second class are precipitated by solutions 
containing a bivalent anion, such as potassium sulfate, K.SO, 
(see p. 261). 

Colloidal silicic acid, which is present in arable soils, retains 
cations in a form which cannot be washed out by the rain, 


x WATER 


it 





354 A TEXTBOOK OF CHEMISTRY 


probably owing to the relations which have just been given, and 
potassium seems to be held much more strongly than magnesium, 
calcium or sodium. ‘This is doubtless of great importance in its 
relation to the fertility of the soil, since potassium is one of the 
most essential elements for the growth of crops. 

If hydrochloric acid is added to the solution of sodium silicate 
drop by drop, instead of suddenly, the silicic acid separates as a 
gelatinous precipitate. If this is dried on the water bath or at a 
slightly higher temperature, or if the colloidal solution obtained 
by the sudden acidification is evaporated and the residue dried, 
the silicic acid becomes almost completely insoluble in acid 
solutions. Such a process is much used in the quantitative 
analysis of silicates. 

The gelatinous precipitate obtained by precipitation has very 
nearly the composition Si(OH),4 (Norton, J. Am. Chem. Soc. 
19, 882 (1897)), but there is considerable doubt whether it is a 
definite compound, since it loses water very easily and its vapor 
pressure is scarcely, if at all, different from that of pure water. 
If dried, however, it loses the last portions of water with difficulty 
and only when heated to bright redness. Opal must be con- 
sidered as a mixture of silicic acids, (SiO2.H2O), but is variable 
in composition, and it cannot be said that the existence of any 
definite compound having the composition of a silicic acid has 
been established. Silicon dioxide, SiOs, resembles carbon 
dioxide, COs, in this respect, but with the difference that while 
carbonic acid, H2COs3, dissociates directly into carbon dioxide 
and water the hydrates of silicon dioxide lose water gradually, 
forming hydrates containing less and less water, that are united 
together to form complex molecules, perhaps somewhat as 
follows : 


o=si/: OH _H{ ON gi—o0 -> H—O0—Si0O—O—SiO—OH 
\OH H—0” 
The final product of the dehydration of the silicic acid would, 
according to this view, consist, not of simple molecules of silicon. 
dioxide but of complex molecules (SiOz), in which the molecules 


NATURAL SILICATES 355 


are held together through the agency of oxygen. The large 
number of complex silicates which are known furnish a strong 
support for this view. The extremely high melting point and 
boiling point of silica, especially as compared with carbon di- 
oxide, also indicate that its molecule is complex and that it does 
not have the simple formula SiOz. 

Natural Silicates. Carbon forms a great number of acids 
in which the complexity is due to the union of carbon atoms with 
each other and with varying numbers of other atoms and groups, 
and it forms only a single acid for which carbon dioxide is the 

Pye Vee 
anhydride. Such a grouping as O=C—O—C=O seems to 
be extremely unstable. As has just been pointed out, how- 
ever, molecules containing silicon atoms united by oxygen seem 
to be stable, and a great variety of minerals are known which may 
be considered as salts of more or less complex silicic acids, all of 
which are derived from the same anhydride, silica, SiO. The 
hypothetical acids from which these natural silicates are derived 
are best classified according to the number of atoms of silicon 
contained in one molecule of the acid. The first two are given 
names similar to the names of the acids of phosphorus, arsenic 
and antimony: 

Orthosilicic Acid H,SiO, 
Metasilicic Acid H.SiO3 
Disilicie Acid HSi207 
Trisilicie Acids H,Si303 and HsgSizsO10 

The following minerals may be given as illustrations of the salts 

of the above acids. 


Calamine, Zn2SiO4.H2O 


Scie ea Mica, KH2Al3(SiOz)3 
ee eciues| Kaolin Als(ciOacH.0 
H4Si04 


Garnet, Ca3Fe2(SiO4)3 


Amphibole 
Metasilicates; | Hornblende ; CaMg3(SiO3)4 


Derivatives of <4 Asbestos 


H.SiO; Tale or soapstone, Mg3H2(SiOs)4 
Beryl, BesAle(SiO3)¢ 


UNIVERSITY OF ILLINOIS 
LIBRARY-CHEMISTRY¥ 


356 A TEXTBOOK OF CHEMISTRY 


Disilicates ; 
Derivatives of 


H,Si,O,7 Serpentine, HyMg3Si.O, (or Mg3SieO7.2 H2O) 
Trisilicates ; 
Derivatives of ete KAISi30¢ 

H,Si;0, Albite, NaAI1Si3;0¢ 
Derivatives of 


HgSisO19 Meerschaum, HiMg2Si3010 


* Calculation of the Formula of a Mineral. The following 
analysis of a garnet may be taken as typical of the analysis of a 
silicate : 














PER CENT irtericae Ratios 
Silica, SiOz alee tle he ee 40.45 FA ele Ty 21.57eb 2066 
Altmina “AlsOs 9h ie ee 19.67 9.26 
Ferricioxide,, Hess) us ee 4.05 131 11.28 iors 
Chromic oxide, CrsO3; . . . . 2.60 0.81 | 
Ferrous oxide, FeO. . . . . 6.90 153 
Calenm oxide CaOrm .* 2) sar, 5.78 1.65 1150418 1 ores 
Magnesium Oxide, MgO . . .}| 20.79 8.32 | 

100.24 














In this and in similar analyses of minerals and rocks it is 
customary to calculate the results as though the elements were 
present as oxides. This .custom was originally based on 
Lavoisier’s system of nomenclature, according to which salts 
were considered as compounds of oxides of the metals with oxides 
of the nonmetals, and it was continued during the first half of 
the nineteenth century largely because of the older electrochemi- 
cal theory, which seemed to give a satisfactory theoretical basis 
for the old nomenclature. This method of calculation is still 
retained because many of the silicates can be prepared by the 
direct union of the oxides of the metals with silica, but also for 
the practical reason that when a silicate analysis is calculated 


DIALYSIS, SEMIPERMEABLE MEMBRANES 357 


in this manner the failure of the percentages found to add up to 
100 shows at once that there is an error in the analysis or that 
some element has been overlooked. 

In calculating a formula for a mineral it is convenient to make 
use of the principle that the amounts of oxygen in the different 
oxides must be in a simple ratio to each other. Thus if we write 
calcium silicate CaOSiOnz, the ratio of the oxygen in the calcium 
oxide (“lime”) must be to that in the silicon dioxide (silica) 
as 1:2. An examination of the quantities of oxygen in the 
second column of figures above shows, however, no simple 
ratio between the amounts of oxygen in the various oxides. It 
is not till we put together the oxygen of the oxides of ferrous 
iron, calcium and magnesium and that of the oxides of ferric 
iron, chromium and aluminium that we obtain numbers which 
form simple ratios. When we do this we obtain, approximately, 
the ratios R’’O: R2’’’03 : Si0g = 3: 1:3, and may write the gen- 
eral formula of garnet, 3 R’’0O.R2/’’03.3 SiOe or R3’’R’’’(SiO.)s, 
in which R”’ stands for ferrous iron, calcium or magnesium and 
R’”’ stands for ferric iron, chromium or aluminium. 

If this formula is written in the ordinary way, it becomes 
3 CaO.Al,03.3 SiO2 or Ca3Al2(SiO.4)3, and we see that garnet is 
a derivative of orthosilicie acid, H4SiO.. 

It is fair to say, however, that the published analyses of many 
of the complex silicates agree only very roughly with the for- 
mulas which have been assigned to the minerals. The condi- 
tions under which such minerals have been formed in nature have 
evidently favored the formation both of complex, isomorphous 
mixtures and of solid solutions of variable composition. 

Dialysis, Semipermeable Membranes. Colloidal silicic acid 
may be separated from sodium chloride by dialysis, with the 
use of parchment or parchment paper. Graham, who first 
carefully studied phenomena of this kind, distinguished between 
crystalloids and colloids in regard to this effect of animal mem- 
branes. As he used these names they imply that crystalline 
compounds, such as sodium chloride, most salts and ordinary 
acids will pass through the parchment, while silicic acid, albu- 


358 A TEXTBOOK OF CHEMISTRY 


men and other substances which do not crystallize, or which 
erystallize with difficulty, will not pass through. While this dis- 
tinction still has considerable force, in a general way, a fuller 
knowledge of the subject of colloids has led to a definition of 
them (p. 261) which is based on quite other properties than their 
relation to separating membranes, and the term crystalloid is 
now little used. 

A membrane or septum which allows one substance to pass 
through it while it prevents the passage of another is called 
semipermeable. A piece of parchment will allow water or salt in 
solution to pass, but is nearly impervious tosilicic acid or albumen. 
It will allow cane sugar to pass, but more slowly than salt. Ifa 
precipitate of copper ferrocyanide, Cuzl'eCeNe, is formed inside 
of the wall of a porous porcelain cup, by placing a copper sulfate, 
CuSO,, solution within and a solution of potassium ferrocyanide, 
Ky4FeCgNe, on the outside, the gelatinous membrane, formed at 
the point in the wall where the two solutions come together is per- 
meable to water but may be made wholly impervious to cane sugar. 
A thin sheet of India rubber is readily permeable to pyridine but 
nearly impervious to glucose in solution in the pyridine. Metal- 
lic palladium is permeable to hydrogen but impermeable to ni- 
trogen and most other gases. 

The mechanism of the permeability is probably different in 
different cases. Palladium dissolves hydrogen very much as 
water dissolves carbon dioxide, but will give it up again toa 
vacuum or to*°any space containing no hydrogen, although it 
may contain another gas. There is some difference of opinion 
about the permeability of copper ferrocyanide. Some authors 
think the pores of the membrane are so fine that molecules 
of water can pass through them while the larger molecules of 
sugar are stopped. Others think that the membrane dissolves 
water on one side and gives it out on the other, as the palladium 
dissolves hydrogen and gives it up. | 

Osmotic Pressure. The passage of a liquid or solution 
through a membrane in the manner which has been described is 
called osmosis. If a concentrated solution of copper nitrate is 


OSMOTIC PRESSURE 359 


placed in the parchment sack, Fig. 91, it will be seen very soon 
that the solution inside of the sack is at a higher level than the 
water on the outside, indicating a greater pressure on the side 
of the solution and showing that 
some water has passed through 


ror ») 


the membrane into the solution. 
A pressure developed in this i 
manner is called osmotic pres- HH 
sure. The pressures developed il 
with the parchment sack will be a 
small, partly because the salt El 


solution as well as the water 
passes through the membrane, 
which does not form a perfect 
septum, and partly because any 
considerable difference of pres- 
sure between the two sides would 
burst the parchment. 

By the arrangement shown in 
Fig. 92 it is possible to measure 
an osmotic pressure of many 
atmospheres. A membrane of 
copper ferrocyanide, Cu2FeCgNe, 
is first prepared, as described in 
the last paragraph, within the 
walls of the porous porcelain 
cup, z. The tube, m, is filled 
with mercury, leaving air in the 
graduated capillary tube. The 
rest of the apparatus is then 
completely filled with a solution 
of sugar or of some other sub- 
stance which is to be examined, and the cup, 2, is placed in 
distilled water. The volume of air in the capillary tube will 
soon begin to diminish, and the contraction will continue till a 
pressure on the solution is developed which will tust prevent the 





Fig. 92 


360 A TEXTBOOK OF CHEMISTRY 


further passage of water through the membrane; the system is 
then in equilibrium. This hydrostatic pressure on the solution, 
necessary to prevent the passage of the solvent through a semi- 
permeable membrane into a solution is the osmotic pressure of 
the solution (van’t Hoff). Itis evident that if equilibrium is 
reached when the volume of the air in the tube is one half the 
original volume, the pressure of the air must be two atmospheres 
and the osmotic pressure must be the difference between this 
pressure and the pressure of the air on the water outside, which 
would be one atmosphere. If the air in the tube is reduced to 
one third its original volume, the osmotic pressure must be two 
atmospheres. 

The osmotic pressures developed in this manner are very con- 
siderable. A tenth-formular solution of cane sugar, Ci2He2Oi 
(containing 34.2 grams in a liter or 3.42 per cent), will give an 
osmotic pressure at 0° of about 2.24 atmospheres, while a formu- 
lar solution of alcohol, C2H;OH (containing 4.6 per cent by 
weight), would give a pressure of more than 20 atmospheres. The 
pressures are very nearly proportional to the absolute tempera- 
ture and, in dilute solutions, nearly the same as though the dis- 
solved substance (solute) were in the form of a gas in the same 
volume and at the same temperature. 

It will be seen, at once, that a measurement of the osmotic 
pressure may be used to determine the molecular weight of a 
compound exactly as the density of a gas or vapor is used for this 
purpose (p. 94). It has been shown, also, by processes of reason- 
ing which it would carry us too far to give here, that there is a 
necessary connection between the osmotic pressure of a solution 
and its vapor pressure, boiling point and freezing point (van’t 
Hoff). The use of the freezing point of solutions to determine 
molecular weights (p. 112) is intimately connected with these 
relations. 

A mental picture of what may be the cause of osmotic pressure 
can be formed by considering the effect of a septum of palladium 
on hydrogen gas. if a bulb of palladium containing nitrogen 
under atmospheric pressure is placed in an atmosphere of hydro- 





GERMANIUM 361 


gen, also under atmospheric pressure, hydrogen will be absorbed 
by the palladium on the outside and given off on the inside until 
the pressure of the hydrogen on the two sides of the wall of 
palladium is the same. The pressure on the inside of the bulb 
wil] be two atmospheres, one atmosphere due to nitrogen and 
one atmosphere due to hydrogen. This is because nitrogen gas 
is discontinuous and offers no resistance to the escape of mole- 
cules of hydrogen from the surface of the palladium. When 
inside, the nitrogen and hydrogen each exert their normal pres- 
sure on the wall as gases. 

In a similar manner, if we have a dilute cation of sugar on 
one side of a membrane of copper ferrocyanide and pure water 
on the other side, the water will pass through till the pressure 
due to the water is the same on both sides. When equilibrium 
is reached the pressure on the side of the sugar solution will be 
greater than that on the side of the pure water by the amount 
of the pressure due to the sugar. It has been shown experi- 
mentally that this pressure is very nearly the same as though 
the solute existed as a gas in the same space. 

It is clear from what has been said that the kinetic theory 
(p. 58) must apply in a modified form to liquids as well as gases. 
We shall find later (p. 397) that it applies to solids also, as shown 
by the specific heats of the elements. 

* Germanium, Ge, 72.6. In 1886 Clemens Winkler found, 
after repeated analyses of a mineral called argyrodite, that the 
sum of the elements found was always 6 to 7 per cent less than 
100. This led him to a further careful study of the material 
and to the discovery of germanium, an element intermediate 
in its properties between silicon and tin, but on the whole resem- 
bling the latter much more closely than the former. It is a 
brittle, grayish white metal, which melts at 958°. It forms the 
compounds GeO, GeS and GeCl, in which it is bivalent and 
GeH,u, GeHCls, GeCli, GeOo, GeS. and K2GeFs, all of which 
correspond to similar compounds of quadrivalent silicon. 

Argyrodite has approximately the composition 4 Ag2S.GeSo. 

Tin, Sn, 119, and Lead, Pb, 207.1, are very distinctly metallic 


362 A TEXTBOOK OF CHEMISTRY 


in their properties in the free state and also in the formation of 
salts in which they are the metallic elements. They give, how- 
ever, the oxides SnOz, stannic oxide, which is the anhydride of 
stannic (H2SnO3), and metastannic acids, and PbOx, lead dioxide, 
which is the anhydride of plumbic acid, H4PbO.. These metals 
and their compounds will be considered more in detail later. 

The elements of the fourth group of the periodic system, 
which are found in the alternate rows, show a closer resemblance 
to silicon and germanium than is shown by tin and lead. It 
seems, on the whole, therefore, better to give an account of them 
here rather than among the metallic elements. 

* Titanium, Ti, 47.9. This element isfound in the mineral rutile 
TiO», in many iron ores, especially in the magnetic iron ore, and 
in small amounts in almost all rocks. Ordinary soils contain, on 
the average, more than half a per cent of titanium oxide and it is 
probably a constituent of the ash of most plants and animals. 
The free element has apparently never been obtained pure, 
largely because of its strong affinity for nitrogen and carbon. 
The purest specimens which have been obtained have a very high 
melting point, 1790°, suggesting the possibility of using the ele- 
ment for the filaments of electric lamps, but no one has succeeded 
in applying it to this purpose. Titanium forms compounds 
in which the element is bivalent (TiCls, TiO, TiS), trivalent 
(TiCl;, Tie03, TioSs, Tis(SO4)3, TiN) and quadrivalent (TiCl,, 
TiOs, TiSe, K2TiFs). 

Titanium tetrafluoride, TiF4, is a white powder which boils 
at 284°, but it is much more easily decomposed by water than 
silicon tetrafluoride so that on treatment of a mixture of silicon 
dioxide, SiO», and titanium dioxide, TiQ., with a mixture of 
hydrofluoric and sulfuric acids, the silicon may be expelled as 
the tetrafluoride on evaporation, while the titanium remains be- 
hind as the oxide. Acid solutions containing titanium give a 
deep yellow color with hydrogen peroxide, due to the formation 
of pertitanic acid. ‘This is used for the detection and estimation 
of titanium and may also be used for the detection of hydrogen 
peroxide. Hydrofluoric acid interferes with the reaction. 


CERIUM 363 


Some compounds of titanium are used as mordants in dyeing 
wool, cotton and leather. 

* Zirconium, Zr, 91, is found most frequently in the mineral 
zircon, ZrSiO4. It is also found in considerable quantities in 
Brazil in the form of the dioxide, ZrO.. The element exists 
in two forms. ‘The amorphous form is a black powder, resem- 
bling carbon. The crystalline form is brittle, melts at about 1700° 
and has a specific gravity of 6.4. Zirconium hydride, ZrHg, is 
a black powder. The chloride, ZrCl4, is hydrolyzed by water. 
Zirconates, which resemble some of the silicates, are prepared by 
fusing zircon dioxide, ZrOe, with metallic oxides, carbonates or 
chlorides. Crystalline silicozirconates have also been prepared. 

Zirconium dioxide glows very intensely in the oxyhydrogen 
flame, giving an even better light than lime, and the light is used 
for intensive illumination in spectroscopy and microphotography. 
The oxide is used with yttrium oxide as the chief constituent of 
the Nernst lamp. The mass conducts electricity well only at 
high temperatures and must be heated in some way to start 
the passage of the current. It was at one time used as the incan- 
descent material for the Welsbach light, but that use has been 
abandoned. There seems to be some possibility of using zircon 
or some of its compounds in the filaments of electric lamps. 

Fused zirconium dioxide resembles fused quartz and cru- 
cibles for laboratory use have been made from it. 

Cerium, Ce, 140.25. In Brazil and in North and South 
Carolina there are found large deposits of a heavy sand, called 
monazite sand, consisting of a complex mixture of heavy, 
insoluble minerals which have been sorted out from other 
materials, partly by the disintegration of rocks in which they 
were originally disseminated, partly by a process of washing 
away the lighter minerals in operations which doubtless occurred 
during very long periods of geological time. These sands con- 
tain a great variety of minerals and are of considerable commer- 
cial value because of the thorium and cerium in them. The ce- 
rium seems to be mostly present as cerium phosphate, CePQs.. 

Metallic cerium is an iron-gray, malleable metal, which burns 


364 A TEXTBOOK OF CHEMISTRY 


easily to cerium dioxide, CeQOs, in the air. Its specific gravity 
is 6.73. An alloy with iron throws off incandescent, burning 
particles on striking with a piece of steel. These will ignite gas 
or alcohol and the alloy is used for gas lighters and for similar 
purposes. 

* Cerium forms three oxides, CesO3, CeOo and CeO3. The 
dioxide, CeQ., forms 1 per cent of the material used for the 
Welsbach mantles. (See below.) Cerous sulfate, Ce2(SOx)3, 
forms a difficultly soluble double salt with sodium sulfate, 
Ce2(So4)3. NagSO4.2 HeO, which is used in separating cerium 
from other minerals. The ceric sulfate, Ce(SOx.)o.4 HO, in 
which the cerium is quadrivalent is readily hydrolyzed to basic 
sulfates. Cerium dioxide places it here; the trioxide places. it 
with the rare earths (p. 502). 

* Hafnium, 178.3, was discovered by Hevesy in 1922. He 
was led to the discovery by Bohr’s theory of the arrangement of 
electrons in successive layers about the nucleus of an atom. 
According to this theory, the element of atomic number 72, 
following lutecium, should no longer belong to the rare earth 
group but should resemble zirconium. Some years earlier 
Moseley had shown that the X-ray spectrum of an element 
contains characteristic lines located in such a manner as to 
fix the atomic number of the element. By applying this 
method to a sample of the mineral zircon from Norway, 
the spectral lines for the element of atomic number 72 
were found. With this clue as a starting point, hafnium was 
finally separated and its properties were found to be those of 
a quadrivalent element closely resembling zirconium. It is 
associated with zirconium in all minerals containing that 
element. 

Thorium, Th, 232.15, is found in monazite sand, probably as 
the dioxide, ThOs, and this forms the practical source for thorium 
compounds, which have become very important for use in Wels- 
bach mantles. Thorium is also found in the mineral thorianite, 
an isomorphous mixture of thorium dioxide, ThOs, and uranium 
oxide, UO:. Thorianite is found in Ceylon. The mineral has 


BORON 365 


proved of unusual interest as the source of radiothorium, one of 
the strongly radioactive elements. 

Apparently pure metallic thorium has not been prepared. 
From the properties of the impure metal it probably has a spe- 
cific gravity of about 12.2 and melts at 1700° or above. 

Thorium forms many salts in which it is quadrivalent, of which 
the sulfate, Th(SO.)o.9 H.O, and the nitrate, Th(NO3)4.12 H.O, 
may be considered as typical. 

Welsbach Mantles. Compounds of thorium and cerium are 
now important articles of commerce for the manufacture of the 
mantles used for the Welsbach light. In the preparation of 
these mantles a web of cotton or, more recently and better, of 
artificial silk is dipped in a solution containing nitrates of cerium 
and thorium. Careful experiments have shown that the best 
results are obtained with a mantle containing 99 per cent of 
thorium dioxide, ThOs:, and 1 per cent of cerium dioxide, CeOr. 
After drying, the cotton or artificial silk is burned out and the 
skeleton of oxides is dipped in collodion and dried. The elastic 
film of nitrocellulose left by the evaporation of the collodion 
holds the fragile skeleton of oxides together for transportation. 

A study of the spectra of the light from mantles made of pure 
thorium oxide, pure cerium oxide and of mixtures of the two has 
shown that thorium oxide radiates comparatively little light of 
the wave lengths of the visible spectrum, apparently because it 
is nearly transparent. Cerium oxide, on the other hand, 
radiates so much energy of the wave length of the ultra red rays 
that it lowers the temperature of the flame too far to produce a. 
satisfactory light. In the mixture, the thorium dioxide, which 
has a very low specific heat and slight power of emission, is 
heated to a high temperature (1500°-1600°), and at this temper- 
ature the cerium oxide has a greatly increased power of radiation 
for rays of the wave lengths of the visible spectrum. Small 
amounts of chromium, platinum, manganese or uranium may 
produce effects similar to those produced by the cerium, but 
none of these give as successful a mixture as that of the oxides 
of thorium and cerium. 


366 A TEXTBOOK OF CHEMISTRY 


Boron, B, 10.82. In some places in Tuscany, in northern 
Italy, steam which is charged with boric acid, H3BQs3, escapes 
from subterranean sources. The boric acid is obtained by bring- 
ing the steam into contact with water, which absorbs the acid 
and by evaporating the solution till it crystallizes on cooling. 
Borax, NagB,O0;.10 H2O, is found in many natural waters and 
especially in Borax Lake, in California. <A calcium salt, called 
colemanite, CagBgQOu.5 H2O, is also found in veins in California 
and is used for the manufacture of borax. 

Preparation, Properties. Amorphous boron may be prepared 
by heating the trioxide, BeO3, or dehydrated borax, NagB.4O;, 
with magnesium. It forms a brownish gray powder. Boron 
may be obtained in a crystalline form by crystallizing it from 
metallic aluminium. The crystals are extremely hard, approach- 
ing the hardness of the diamond. Boron melts at 2300°. 

Boron Trioxide, B.O3, Borax Beads. ‘The trioxide or boric 
anhydride, B.Os, is formed by heating either of the oxygen acids 
of boron to a high temperature. It is volatile only at a white 
heat and on this account it will decompose the salts of almost 
all other acids, when heated with them to a high temperature. 
It does this in spite of the fact that in solution boric acid is one 
of the weakest acids known — so weak, in fact, that its solution 
does not taste sour and does not redden litmus. When fused 
with salts or with the oxides of metals, boric anhydride combines 
with the oxides to form borates, just as silicon dioxide combines 
with oxides to form silicates. Borax glass, NagB,O7, which may 
be considered as boric anhydride that has been only partly 
neutralized by sodium oxide (Na,O.2 B.O3), decomposes salts 
or combines with oxides of metals in the same manner, and 
almost all of the borates of metals prepared in this way remain 
dissolved in the anhydride or combined with it, giving a mixture 
of borates, which solidifies to a clear glass on cooling. Many 
of these glasses have characteristic colors and so the borax bead 
is much used for the detection of metals in mineralogy and 
qualitative analysis. 

Some of the colorless borates may be fused with silicates to 


OTHER COMPOUNDS OF BORON 367 


form clear, transparent glasses, which are much less soluble than 
ordinary glass. Such borosilicates are now used in the manu- 
facture of glassware for chemical laboratories. 

Boric Acid, H3;BO;3. If sulfuric acid or hydrochloric acid is 
added to a moderately concentrated solution of borax, 
NaeB4O7.10 H,O, in an amount equivalent to the sodium 
present, orthoboric acid, H3BQOs3, crystallizes from the solution 
on cooling. It is usually called simply boric acid and is a very 
weak acid, having no sour taste and none of the corrosive 
qualities of ordinary acids. It dissolves in about 25 parts of 
water at 15°. The solution has powerful germicidal properties 
and is often used for an eye wash. Both boric acid and borax 
have been used as food preservatives, but such a use is con- 
demned because the substances produce poisonous effects when 
taken in considerable quantities. 

Other Acids of Boron. At 100° boric acid loses one molecule 
of water and is changed to metaboric acid, HBOs, and at 140° 
it loses still more water and is converted into pyroboric acid, 
H,B,0O;. 

Borax, Na2B,O07.10 HO, is usually considered as a salt of 
pyroboric acid, H.B,O;, but it may with equal propriety be 
called an acid salt of boric acid and the formula written 
NazHio(BO3)4.5 HeO. When heated it loses water and finally 
melts to a clear glass consisting of sodium pyroborate, Na2B4O7. 
The conduct of this glass toward metallic oxides and its use in 
blowpipe analysis have been mentioned above. Because of 
these properties it is used as a flux in assaying and in metallurgy. 
It is sometimes sprinkled over iron which is to be welded. In 
this use it dissolves the oxide of iron on the surface, forming an 
easily fusible glass which is pushed out from between the two 
surfaces when the joint is hammered, leaving clean surfaces of 
iron to unite in the weld. 

Borax is used for laundry purposes, having the properties of a 
milk alkali. Its use as an antiseptic in milk and other foods is 
forbidden in most countries. 

Sodium Perborate, NaBO3;.4 H.O, has recently come into use 


368 A TEXTBOOK OF CHEMISTRY 


in laundries because it combines the properties of a mild alkali 
with the bleaching properties of hydrogen peroxide. Its struc- 
ture, from the methods of preparation, must be closely related 
to that of hydrogen peroxide, H—O—O—H, and is probably 
Na—O—O—B = O. 

Other Compounds of Boron. Boron forms a nitride, BN, a 
chloride, BCl;, and a sulfide, BeSs3, all of which are hydrolyzed 
by water. The fluoride, BF3, gives boric acid and fluoboric 
acid by hydrolysis, resembling silicon fluoride, SiF4, in this 
respect: 

BF; + 3 HOH = H;BO; + 3 HF 
HF + BF; = HBF, 


Fluoboric 
Acid 


If boric acid or a borate is warmed with alcohol and con- 
centrated sulfuric acid, an ester of boric acid, (CsHs)3BOs, is 
formed. This is volatile and gives a green color to the alcohol 
flame, as it burns. 

If a piece of turmeric paper is dipped in a solution of boric 
acid or of a borate which had been made faintly acid with hydro- 
chloric acid, the paper assumes a very characteristic red color 
on drying. 

EXERCISES 

1. How many grams of magnesium should be mixed with 3 grams of 
sand to form silicon? How many grams to form magnesium silicide ? 

2. In the first case above how many cubic centimeters of hydro- 
chloric acid, sp. gr. 1.10, containing 20 per cent of the acid will be 
required to dissolve the magnesium oxide formed ? 

3. What is the formula of a mineral having the following com- 
position ? 


Per Cent 
Sila, ios ee ena. ea ae eee 66.21 
Alumina, “A lOs sreece 2 oy hare a ee eee 19.16 
Potassium oxide kel). el”) 2) eee eee 7.38 
Soditimmoxides NaO ny tet) Shaiya 1.25 


CHAPTER XXII 


METALLIC ELEMENTS. DIFFERENCES BETWEEN METALS 
AND NON-METALS. PREPARATION OF COMPOUNDS 


Metals and Non-metals. The general, physical differences 
between the metals and non-metals are familiar to every one. 
Iron, copper and gold may be taken as typical of the first class, 
sulfur, phosphorus and carbon of the second. The metals are 
opaque, except in excessively thin layers, are malleable and ductile 
and good conductors of heat and of electricity. The non-metals 
are transparent or translucent, brittle, poor conductors of heat 
and nonconductors or poor conductors of electricity. It is 
true that these differences, which are so marked in the elements 
spoken of above as typical of the two classes, are not equally 
marked in all cases and that there is a gradation in these proper- 
ties such that some elements stand in a borderland between the 
two, but the physical properties mentioned are clearly those of 
metals on the one hand and of non-metals on the other. 

In chemical properties the metals and non-metals show equally 
strong contrasts. In these properties the most typical metals 
may be considered as sodium and potassium and the most 
typical non-metals as fluorine and chlorine. The chlorides and 
even the hydroxides of the former ionize in solution in such a 
manner that the metal becomes the positive ion. The chlorides 
of the nonmetallic elements, on the other hand, are hydrolyzed 
by water with the formation of hydrochloric acid and an acid 
containing the non-metal. The compounds of the non-metals 
with oxygen and hydrogen ionize in solution with the formation 
of hydrogen ions and some complex group containing the non- 
metallic element. As is the case with the physical properties, 
however, the chemical properties of the elements show continu- 


ous gradations from those of the metals to those of the non- 
369 


370 : A TEXTBOOK OF CHEMISTRY 


metals. Some of the chlorides of the metals are hydrolyzed by 
water and some halides of elements which are nonmetallic in 
most of their properties may be formed in the presence of water. 
Further than this, some elements as they combine with more and 
more oxygen may pass from metallic to distinctly nonmetallic 
properties, while nitrogen or even sulfur, when combined with 
hydrogen or hydrocarbon radicals, may form radicals, such as 
NH; or (CH3)3S, whose hydroxides are bases. 

The most typical of the metallic elements are univalent toward 
chlorine and toward oxygen and form no compounds of any kind 
in which it is certain that they have a higher valence. The most 
typical nonmetallic elements on the other hand show a varying 
valence, especially toward oxygen. 

The development of the electron theory makes it seem possible, 
or perhaps we may even say, probable, that both the physical 
and chemical properties which constitute the differences between 
metals and non-metals are very largely occasioned by the dif- 
ferences in the conduct of the atoms of the elements toward 
electrons. The passage of electricity through a metallic con- 
ductor consists in the flow of a stream of electrons, and metals 
are supposed to be good conductors because the electrons pass 
easily from one atom to another throughout the mass. For 
almost the same reason the atom of the metal becomes a posi- 
tive ion in solution because it easily loses one or more electrons 
— the electrons being units of negative electricity. Metals are 
also good conductors of heat because the rapidly moving electrons 
transfer energy from one atom to another throughout the mass, 
and they are opaque because the electrons and atoms absorb 
and diffuse the light vibrations falling on the surface. 

Classification of the Metals. As in the case of the non-metals, 
the periodic system furnishes the most satisfactory classification 
for the metals. The relationships brought out in this way are not 
always so close as might be desired, and there are even sharper 
contrasts between the alternate metals of the first group than 
those which have been noticed among the elements of the sixth 
and seventh groups. 


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METALLIC ELEMENTS 


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dOQOUD OUI 


ore A TEXTBOOK OF CHEMISTRY 


All of the elements sufficiently well known so that the atomic 
weights are given in the International Table of Atomic Weights 
are included in the table, with the exception of hydrogen. The 
nonmetallic elements are inclosed in brackets. 

An inspection of the table shows that even if the noble gases 
are counted as nonmetallic elements more than three fourths 
of the elements known are metallic in character. In spite of this, 
the chemistry of the metallic elements offers less variety than that 
of the non-metals — largely because the atoms of the metals 
show less tendency to combine with each other, or indeed with 
other elements to form complex radicals. Metals usually sep- 
arate by themselves as positive ions in solution, while the non- 
metals more often form complex ions, such as NO3, SO, or ClO3. 

Melting Points of the Elements. The following table gives 
all melting points of the elements which are known with some 
degree of accuracy. The melting points of the elements whose 
names are printed in capitals have been very carefully deter- 
mined, and are used as standard temperatures for calibrating 
thermometers and pyrometers. As nearly as may be, all values, 
in particular the standard points, have been reduced to a common 
scale, the thermodynamic scale. 

Preparation of Chemical Compounds. It is clear from many 
illustrations which have been given in the preceeding pages that 
many chemical reactions result in an equilibrium between . 
reacting compounds such that action seems to cease only because 
two opposing reactions proceed with equal velocity in opposite 
directions. From a theoretical standpoint we are probably 
justified in considering all reactions as reversible, but in some of 
them the velocities of the reactions in opposite directions differ 
so greatly that the equilibrium lies very far on one side and the 
reaction is practically complete. ‘Thus in the very simple case 
of the reaction between hydrogen and oxygen: 


2 H, + O. = 2 H.O 


at 1000° when equilibrium is reached, only 3 parts in 10,000,000 
of the oxygen and hydrogen will remain in the form of the free 


373 


MELTING POINTS OF THE ELEMENTS 


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374 A TEXTBOOK OF CHEMISTRY 


gases (p. 61). At ordinary temperatures the velocity of com- 
bination and dissociation becomes so slow, in this case, that 
either a mixture of oxygen and hydrogen or water may remain 
apparently unchanged for an indefinite length of time, but we 
know that such a condition of apparent equilibrium, dependent 
on the slow velocity of a reaction, is a state of instability. Such 
conditions are, however, very common, and it has been pointed 
out that the varying velocity of different possible reactions is 
of very great practical importance in organic compounds. 

A very large number of the compounds of the metals are 
electrolytes, and many of these can be prepared in aqueous solu- 
tions. The velocity of ionization reactions is so great for all 
ordinary electrolytes and the interaction between ions in solu- 
tion is so rapid that equilibrium is reached almost instantane- 
ously. The amount of each substance present in this equilib- 
rium is often influenced by some property of one of the substances 
which removes it from interaction with the others. Such prop- 
erties are, especially, volatility, solubility, degree of ionization 
and the formation of complex ions. 

Effect of Volatility. If concentrated sulfuric acid is dropped 
into a concentrated solution of salt, there is, at first, no apparent 
action. In the solution the following reactions will be almost 
instantly in a state of equilibrium : 


NaCl = Nat + Cl- 
H.SO, = Ht + HSO,- 
His Clas aH Gl 
Nat + HSO. = NaHSO, 
It is to be noticed that the reaction as usually given 


NaCl + H.SO, = HCl + NaHSO, 


is merely the final result of intermediate ionization reactions and 
probably occurs, at first, in very insignificant amounts. 

As more sulfuric acid is added, a point will be reached when 
the solution becomes saturated with hydrochloric acid. The 
addition of a further amount of sulfuric acid must now result 


VOLATILE COMPOUNDS 31D 


in the escape of hydrochloric acid in the gaseous form, as mole- 
cules of the compound, HCl. ‘This will disturb the equilibrium 
of the reaction: 
Elie tet SL aeec ey BL] 

in such a manner that more of the hydrogen and chloride ions 
will unite to form un-ionized hydrochloric acid, HCl. This, in 
turn, will cause the formation of an increased number of sodium 
and chloride ions from the salt, NaCl, and of hydrogen and hy- 
drosulfate ions, HSO,-, from the sulfuric acid. Of course, the 
more concentrated the original solution of salt is, the more hy- 
drochloric acid can be obtained in gaseous form, and if we start 
with a saturated solution and a large excess of solid salt, the latter 
will pass into solution as the reaction proceeds and nearly all of 
the chlorine may escape as hydrochloric acid. This result 
occurs because the hydrochloric acid is a gas and escapes from the 
mixture as a gaseous phase and in spite of the fact that the 
ionization reactions of hydrochloric and of sulfuric acids result 
in the formation of more of the sulfuric than of the hydrochloric 
acid when the two acids are present in equivalent amounts. 

When a solution of hydrochloric acid is dropped into a solu- 
tion of sodium carbonate the following reactions occur : 


HCl 3H s Cle 
NazCO3 = Na* + Nat + CO3= 
CO;s- + Ht 2 HCO;7 
Ht + HCO;- == H:CO; 
H.CO; = HO + CO, 
Nat + Cl- 2 NaCl 
Nat + HCO; == NaHCO; 

In the third of these reactions the equilibrium is very far in- 
deed toward the formation of the hydrocarbonate ion, HCO;", 
and very few hydrogen ions, H*, can remain in the solution so 
long as any carbonate ions are present. Since many more hydro- 


genions separate from carbonic acid, HeCOs3, when that is present, 
than separate from the hydrocarbonate ion, HCO37, carbonate 


376 A TEXTBOOK OF CHEMISTRY 


ions, CO;-, and carbonic acid, HeCOs, cannot exist to any large 
extent in the same solution. It will be seen from the above 
that carbonic acid cannot be formed in the solution in suffi- 
cient amount for the rapid escape of carbon dioxide! until all of 
the carbonate ions present have been converted to hydrocarbon- 
~ate ions. In other words, enough hydrochloric acid must be 
added to complete the reaction : 


Na2CO3 + HCl = NaHCO; + NaCl 


before carbonic acid can be formed in sufficient amount for the 
rapid escape of carbon dioxide. The addition of a small addi- 
tional amount of acid, however, causes the fourth and fifth 
reactions to occur and the escape of carbon dioxide begins. 
From this point on the gas will escape in almost exact proportion 
to the acid added, since the equilibrium is far to the right in the 
first, fourth and fifth reactions. 

It is clear from the above discussion that the formation of 
a volatile product, which escapes as a gaseous phase, has the 
same effect on the course of a reaction when it is a secondary 
product, formed by dissociation, as when it is formed directly 
by the ionization reactions, as was the case with hydrochloric 
acid. 

Effect of Insolubility. When a compound separates as a solid 
phase from a solution, the effect upon the equilibrium is exactly 
analogous to the effect of volatility. Thus in the reactions which 
occur on mixing a solution of salt with a solution of silver nitrate: 


NaCl = Nat + Cl- 
AgNO; = Agt + NO; 
Agt + Cl” = AgCl 
Nat + NO3;” == NaNOs 


1A solution of sodium bicarbonate, NaHCO:;, contains some car- 
bonie acid, H,CO;, formed by the reaction : 


Nat+ HCO3;- + Ht + OH--= H2CO3 + Nat++ OH- . 


Such a solution will lose carbon dioxide on boiling or on exposure 
to the air. 


SOLUBILITY PRODUCT syle 


the fact that silver chloride is only very slightly soluble causes 
it to separate from the solution as a precipitate and shifts the 
equilibrium of the first three reactions toward the formation 
of thiscompound. The removal of a substance from the mixture 
as a solid phase has exactly the same effect upon the various 
equilibria involved as the removal of a volatile product. 

Effect of a Common Ion. Solubility Product. When a sub- 
stance having one of the ions of a difficultly soluble uni-uni-. 
valent salt, that is, a salt such as silver chloride, AgCl, in which 
both the metallic and acid radicals are univalent, is added to a 
saturated solution of the salt, some of the difficultly soluble salt 
will usually be precipitated. Thus the addition of a few drops 
of a solution of silver nitrate to a saturated solution of silver 
chloride will cause the separation of some silver chloride from 
the solution and the addition of hydrochloric acid will also 
cause precipitation in a saturated solution of silver chloride. 

The precipitation depends on the following reactions : 


AgCl => Agt + Cl- 
HC] 2 H+--.Cl- 


In the first reaction the quantities of the three substances, 
silver chloride, AgCl, silver ion, Agt, and chloride ion, Cl~, must 
be present in definite amounts in a saturated solution of silver 
chloride at a given temperature. At 18° the amount of silver 
chloride in the three forms dissolved by a liter of water is 1.6 
milligrams of which about 0.4 of a milligram is chlorine. IH, 
now, hydrochloric acid is added, the number of chloride ions 
must be largely increased, and this will cause the formation of 
more un-ionized silver chloride. As the solution is already 
saturated with that compound, the silver chloride formed will 

be precipitated. 

Under such conditions as these, when a solid, difficultly 
soluble uni-univalent salt is in equilibrium with a solution con- 
taining a slight excess of one of its ions, the equation: 


Cagt X Ca- = a constant 


378 A TEXTBOOK OF CHEMISTRY 


has been found to be true. Cy,,+ and Cq- are the concentra- 
tions of the ions of the salt in the solution.! 

If, for instance, we add enough hydrochloric acid so that the 
solution is tooo Normal for hydrochloric acid, one liter will con- 
tain 35.5 + 0.4 = 35.9 milligrams of chloride ions, Cl, or nearly — 
one hundred times as much as before, since at these dilutions 
both the silver chloride and hydrochloric acid will be almost 
completely ionized. Under these conditions, in order that the 
product Ca,+ X Ccj- may remain constant the quantity of the 
silver ions must become 7% as large as before. Since the silver 
chloride is almost completely ionized at these concentrations, it 
will be seen that the addition of this small quantity of hydro- 
chloric acid will reduce the quantity of silver in the solution to 
about one hundredth of the original amount. The importance 
of these relations for the precipitation of difficultly soluble salts 
in quantitative analysis will be easily seen. This principle is 
often called the constancy of the solubility product, and may be 
stated as follows: In any dilute aqueous solution saturated with 
a slightly soluble uni-univalent salt the product of the concentrations 
of the ions of this salt 1s constant at a given temperature. 

In some cases the addition of a salt having a common ion will 
cause precipitation in a saturated solution containing a uni- 
bivalent salt, such as lead chloride, PbCls, or in one containing 
a bi-bivalent salt, such as calcium oxalate, CaC.Oux, but there are 
many other cases in which such precipitation does not occur. 
No general rules governing the conduct of such salts can be given 
in the present state of our knowledge. In some cases this failure 
of precipitation may be due to the formation of intermediate 
ions, such as PbCl~. In others it is probably caused by the for- 
mation of more or less stable complex ions. 

Formation of Complex Ions. If solutions of silver nitrate, 
AgNOs, and potassium cyanide, KCN, are mixed in equivalent 


1 This relation can be derived theoretically from the ionization 
reaction : Ascl== ueMexine 


if it is assumed that in very dilute solutions the osmotic pressures 
of the ions are proportional to their concentrations. 


COMPLEX IONS. IONIZATION 379 


proportions, nearly all of the silver ions and nearly all of the 
cyanide ions will be removed from the solution as a precipitate, 
in accordance with the equation : 
AgNO; + KCN = AgCN + KNO; 
Silver 
Cyanide 

If, however, more of the potassium cyanide solution is added, 
the silver cyanide will dissolve. This seems to be in direct con- 
tradiction to the principle of the solubility product given above. 
A more careful examination shows that the apparent contradic- 
tion is due to the fact that the solution no longer contains an 
appreciable number of silver ions, Ag*. This can be shown in 
three ways: 

1. Sodium chloride, NaCl, will cause no precipitate of silver 
chloride, AgCl, to form in the solution. 

2. If an electric current is passed through the solution, the 
silver in the solution is carried toward the anode, not toward the 
cathode. This shows that the silver atoms form part of complex 
ions which carry negative charges instead of existing in the solu- 
tion as silverions. If the silver were in the form of the ion, Ag”, 
the electric current would carry it toward the cathode, or negative 
pole. 

3. By evaporating the solution a definite compound, potassium 
silver cyanide, KAgC.N2, may be crystallized from it. ‘Trans- 
ference experiments have shown that the ions of this compound 
are K* and AgC2Ne-. (See p. 320.) 

Many other complex ions are known which are very different 
in their properties from the ions which unite to form them. 

It is clear from this illustration that the principle of the solu- 
bility product depends on the character of the ions which are 
actually present and not on the amount of a given element which 
may be present in the solution. 

Degree of Ionization. If two electrodes in circuit with a 
battery and ammeter are placed at the ends of a narrow, rectangu- 
lar cell filled with distilled water (Fig. 93) and a concentrated 
solution of potassium chloride is put in the bottom of the cell, 


380 A TEXTBOOK OF CHEMISTRY 


the ammeter will indicate the passage of an electric current. 
If, now, the solution is stirred so that the potassium chloride is 
uniformly mixed with 
the water, it will be 
seen that the current 
increases. Since the 
amount of potassium 
chloride in the cell is 
not changed by mixing 
the solution with the water above it, it is evident that a given 





Fig. 93 


amount of the salt is more effective in conveying the current in 
a dilute solution than in a concentrated one. This is explained 
by the kinetic theory and theory of ionization by supposing, 
first, that only the ions, K* and Cl’, take part in the conduc- 
tivity of the solution, and second, that in the ionization reaction : 


Kee kore ie 


the equilibrium is displaced to the right by dilution because the 
potassium and chloride ions meet each other to unite less fre- 
quently in the dilute solution, while the tendency to separate 
into ions is about the same in one solution as in the other. If 
the conductivity of solutions of potassium chloride at increasing 
dilution is measured in this way, the results given in the follow- 
ing table are obtained. Of course, for the dilute solutions it 
would be necessary to calculate from the conductivity of a 
relatively small quantity of the solution what the conductivity 
of the whole quantity of potassium chloride would be. 


CoNDUCTANCE OF SOLUTIONS oF PoTAsstum CHLORIDE 


A A / 

A 

N Potassium Chloride (74.5 grams in 1 liter ) 98.28 mhos! 0.76 

N/10 Potassium Chloride (74.5 grams in 10 liters) 111.97 mhos 0.865 
N/100 Potassium Chloride (74.5 grams in 100 liters) 122.37 mhos 0.945 
N/1000 Potassium Chloride (74.5 grams in 1000 liters) 127.27 mhos 0.983 
N/10000 Potassium Chloride (74.5 grams in 10000 liters) 129.00 mhos 0.996 
N/oo Potassium Chloride (74.5 grams in o liters) 129.5 mhos_ 1.000 


1 The conductance in reciprocal ohmsis the reciprocal of the resist- 
ance inohms. The unit for conductance is one mho. 


DEGREE OF IONIZATION 381 


Under / is given the conductance in reciprocal ohms! of one 
gram molecule of potassium chloride in the volumes of solution 
stated, when placed between two electrodes one centimeter apart 
and sufficiently large to contain the whole solution between 
them. For a normal solution the electrodes might be 25 * 40 
em., giving a surface of 1000 sq. cm. 

It will be seen that the values for the conductances with in- 
creasing dilution form a converging series from which a value of 
129.5 for infinite dilution can be calculated. In accordance with 
the theory that the electric current is carried only by the ions and 
that at infinite dilution the compound is completely ionized, 
it is easy to calculate the degree of ionization by dividing the 
conductance for any given concentration by the conductance 
at infinite dilution. Thus the fraction of a normal solution of 
potassium chloride in the form of ions is =a = 0.76. The 
values given in the last column of the table above have been 
calculated in this manner. 

In a previous chapter it was pointed out (p. 112) that a solu! 
tion containing 46 grams of alcohol in 10 liters of water freezes 
at — 0.184. One containing 74.5 grams of potassium chloride 
in 10 liters freezes at — 0.343°. It would seem from this that if 
the potassium chloride were completely ionized, the depression 
of the freezing point would be twice that of the alcohol, or 0.368°. 
This would correspond to an increase of 0.184° in the depression 
of the freezing point. Comparing this with the increased de- 
0.159 
0.184 
= (0.864 as the fraction ionized. This is in very close agreement 
with the results found by the conductivity method. 

Cane sugar is hydrolyzed by dilute acids to a mixture of glu- 
cose and fructose (p. 334). It is found that a given amount of 
hydrochloric or nitric acid causes a much more rapid hydrolysis 
than an equivalent amount of acetic acid. If we assume that 
the rate of hydrolysis is proportional to the number of hydrogen 
ions present, it is possible to calculate from series of experiments 





pression which is observed (0.343° — 0.184°) we have 


382 A TEXTBOOK OF CHEMISTRY 


with different acids the relative ionization of the acids. The 
results of such experiments are, again, in general agreement with 
the results obtained by the conductivity and freezing-point 
methods. 

The three methods all indicate that there are very great differ- 
ences in the degree of ionization of different compounds. The 
fact that the results obtained by three methods so radically 
different are in approximate agreement is very strong evidence 
that the three effects have a common basis, and no satisfactory 
theory other than that of ionization has been proposed to account 
for the phenomena observed. The lack of complete agreement 
indicates that some factors which are not yet entirely understood 
modify the effects, just as the mutual attraction of the molecules 
of gases prevent them from obeying exactly the laws of Avo- 
gadro and of Boyle. 

The following table gives the degrees of ionization of a number 
of common substances. In the measurement and calculation 
of these values some factors which have not been discussed are 
involved, but the fundamental principle used is a comparison of — 
the conductivity of the solution for which the degree of ionization 
is given with the conductivity of the same substance when it is 
completely ionized. 


Fractions I[ONIzED 


As the mobilities of the Htand OH™ ions change with the ion con- 
centrations, the conductance ratio (A/ A ) does not correctly represent 
the fractions ionized for strong acids and bases. Recent investigations 
have shown that these substances are ionized to about the same extent 
as salts of the same type, and values for them are so given in the table. 
With salts of acids and bases having bivalent ions, the ionization rela- 
tions are complicated by the presence of intermediate and complex ions, 
for instance HgCl. gives besides Hgt* and Cl-, the ions HgCl~ and 
HgCl,-. The figures given must therefore be regarded only as relative 
measures of the tendencies of these substances to form ions. 


DEGREE OF IONIZATION 383 














TENTH HUNDREDTH 
SupstaNon tons Sen leat 
Wes () Ve 00) 
INitriG ACids a. oo. EN CO) ome 0.86 0.94 
Hydrochloric Acid . . HCl 0.86 0.94 
Hydrobromic Acid . . Fide bree 0.86 0.94 
Hydroiodic Acid . . pig beak 0.86 0.94 
Chloric Acid . . . . H+, GlO3;> 0.86 0.94 
Perchloric Acid . . . Hie Clois 0.86 0.94 
Permanganic Acid . . H+, MnQ,- 0.86 0.94 
Tartaric Acid .. . H-, HC,H.0,— 0.098 0.31 
Acetic Acid Se) oe ae ‘Ee. C.H30.- 0.013 0.043 
Hydrocyanic Acid . . Hz.CN] 0.00008 0.00026 
DOr Acid 2°. a: H+, H.BO3- 0.00008 0.00026 
Hydrosulfuric Acid. . HP ahos 0.00095 0.00031 
Sulfurous Acid . . . Hz HSO,- 0.50 0.70 
Carbonic Acid .. . HA HCO:= 0.0017 0.0055 
Phosphoric Acid . . H+, H,PO.- 0.28 0.64 
NCUOl an, elena tae. cs H+ , C,H;O- 0.00011 0.00036 
pUunuriC Acid, sy oo. aS (iY, Hae SOi—) 0.608 0.832 
Oxalicu cid apie ites aria Ur toatl. COs) (a kOol 7, 0.398 
Sodium Hydroxide. . Nat, OH- 0.86 0.94 
Potassium Hydroxide . Sag MO a= 0.86 0.94 
*Tetramethyl Ammo- 

nium Hydroxide. .| N(CH3)4+, OH- 0.85 0.94 
Barium Hydroxide. . |4(Bat+,OH-,OH~-)| 0.76 0.88 
Ammonium Hydroxide! NH, OH- 0.013 0.042 
VWVateres. chutes | cae Ha.OH- 0.0000001| 0.0000001 








1This gives the fraction of the total ammonia in the solution 
which is in the form of ammonium, NH,'t, and hydroxide, OH-, 
ions. Actually much of the ammonia is present as NH; and the 
proportion of real NH,OH ionized is much larger than that given 
in the table. 

2 The value for water gives the fraction of a mol of water in 
one liter, which is in the form of ions at 25°. For comparison with 
the other values in the table the values under V — 10 must be 
divided by 10 and those under V — 100 must be divided by 100. 
Thus one liter of tenth-normal hydrocyanic acid contains 0.000008 
mol in the form of ions and one liter of hundredth-normal acid 
contains 0.0000026 mol in that form, while one liter of pure water 
contains 0.0000001 mol. 


384 A TEXTBOOK OF CHEMISTRY 


Fractions IonizEp — Continued 











TENTH HUNDREDTH 
Soaavk non Ions FORMULA FORMULA 
SOLUTIONS SOLUTIONS 
Ve— 10 V = 100 

Sodium Chloride . . Nat, Cl- 0.852 0.935 
Potassium Chloride. Be rogh BH Fr | 0.855 0.941 
Ammonium Chloride . WHS Gre 0.852 0.936 
Sodium Nitrate. . . Nat, NO3- 0.832 0.932 
Potassium Nitrate . . Ke NOs 0804 2 120.035 
Silver Nitrate .. . Agt, NO3— 0.816 0.931 
Potassium Chlorate. K+, ClO3- 0.824 0.933 
Sodium Acetate 2. oe Nat, C2H3;02- 0.795 0.914 
Potassium Cyanide. 4 Gs aye C15) ee 0.84 0.93 
Sodium Bicarbonate . Na*, HCO. 0.84 0.93 
Potassium Sulfate . .| 4 (K+, K+, SOs=) | 0.724 | 0.870 
Sodium Sulfate . . .|3(Nat,Nat,SO.=) | 0.704 0.857 
Normal Sodium Car- 

bonaté.«. «. 4. -~eus (Nat, NarsGO.=y a7 1 0.86 
Calcium Sulfate. . .}| 4% (Catt, SQs=) 0.625 
Zane Ghloride= os. Gown | bg hohe Cote BO tad 0.86 
Pane oiiate 1.8 bee 4 (Znt++, SO.=) 0.405 0.633 
Copper Sulfate . .- .| 3 (Cutt, SQ,=) 0.396 0.629 
Mercuric Chloride . .|3 (Hg++, Cl-, Cl-)}| 0.01 0.03 














(This table was prepared by Dr. D. A. MacInnes.) 


Effect of Degree of Ionization. Neutralization. When a 
volatile product escapes from a mixture, or when a solid substance 
is precipitated, the equilibria of the reactions which lead to the 
formation of such compounds are shifted in such a way as to 
promote their formation. From the table which has just been 
given it is evident that certain ions cannot exist in any number in 
solutions which contain certain other ions. There can be very 
few hydrogen ions, H*, in solutions containing hydroxide, OH-, 
hydrosulfide, HS~, or hydrocarbonate, HCO3", ions, or in the 
presence of the ions of any of the weak acids. If the solutions of 
two substances giving hydrogen and hydroxide ions are mixed 
in equivalent amounts, the two ions unite to form water and 


HYDROLYSIS 385 


the resulting solution is neutral, if both of the compounds have 
the same degree of ionization : 


HCl = Ht + Ch 
NaOH = Na* + OH- 
H* + OH’ = H.O 


The equilibrium of the last reaction is so far toward the forma- 
tion of water that both hydrogen and hydroxyl ions and also 
practically all of the hydrochloric acid and sodium hydroxide 
disappear from the solution. This is the ordinary reaction of 
neutralization. A solution is neutral when the number of hy- 
drogen, H*, and hydroxide, OH, ions is equal. 

An illustration of the effect of adding hydrochloric acid to a 
solution containing carbonate ions was given above (p. 375). 

Hydrolysis. It will be seen from the table that the ionization 
of carbonic acid, H2COs, to hydrogen, H*, and hydrocarbonate, 
HCO; , ions is very slight. The further ionization of hydrocar- 
bonate ions to hydrogen, H~, and carbonate, CO;-, ions must 
be almost vanishingly small in the presence of an excess of hy- 
drogen ions. A solution of sodium carbonate, Na2CO3, which 
may ionize, at first, as follows: 


Na:CO; = Na* + Na* + COs 
immediately gives with the ions of water 
CO;- + H* + OH- 2 HCO;- + OH- 


This formation of hydrocarbonate ions, HCO,-, reduces the 
number of hydrogen ions and the solution must react alkaline 
because of excess of hydroxide ions, OH’, present. Similar re- 
actions, due to the ions of the water and called hydrolysis, 
occur with the alkali metal salts of all of the very weak acids 
and especially with the salts of weak dibasic and tribasic acids, 
such as sulfides, borates and phosphates. 

The hydroxides of ferric iron, Fe(OH)s3, aluminium, Al(OH)s, 
chromium, Cr(OH)3;, and of many other elements are so insoluble 
that it is impossible to determine the degree of their ionization, 


386 A TEXTBOOK OF CHEMISTRY 


and there is reason to think that the ionization of the hydroxyl 
in basic compounds such as FeCl.OH, which probably exist in so- 
lution, is very trifling. In a solution of ferric chloride, therefore, 
such reactions as the following occur: 


FeCl; == FeCl.t + Cl- 
FeCl,* + OH- + Ht = FeCl,.OH + H* 


Basic 
Ferric Chloride 


The resulting solution, which contains an excess of hydrogen 
ions, will react acid. | . 

These facts may also be stated in the form: Weak acids cannot 
neutralize strong bases and weak bases cannot neutralize strong 
acids completely, when mixed in equivalent proportions, because 
of the hydrolysis of the salts formed. It is chiefly because of 
these relations that acids and bases which are but slightly ionized 
in their solutions are called “ weak.”” The degree of ionization 
furnishes the only satisfactory basis for classifying acids or 
bases as “ weak” or “ strong.” 

Illustration of the Strength of Acids. If solutions of potassium 
iodide (KI, 6 grams per liter) and potassium bromate (KBrOs, 
1 gram per liter) are mixed in equal proportions, no reaction 
occurs between them, but on the addition of an acid, iodine is 
liberated and colors the solution yellow or brown: 


6 HI + HBrO; = HBr+ 3 I, + 3 H.O 


If the acid is dilute, the reaction is sufficiently slow so that its 
progress can be noted by a slow change in the color. The rate 
of the change is proportional to the number of hydrogen ions 
present. If equal volumes of the iodide-bromate solution are 
placed in three glasses and there is added to these, respectively, 
5 ec. of tenth normal hydrochloric acid (3.65 g. HCl per liter), 
5 ec. of tenth normal oxalic acid (5.38 g. H2C204.2 H.O per 
liter) and 5 ec. of tenth-normal acetic acid (6.0 g. HC2H3O0z2 per 
liter), a very rapid change in color will occur in the first solution, 
a much slower change in the second and a very slow change, 
indeed, in the third. On the other hand, if 5 cc. of each acid 
are placed in three beakers and 5 cc. of tenth-normal sodium hy- 


USE OF INDICATORS 387 


droxide are added to each, the three solutions, containing sodium 
chloride, NaCl, sodium oxalate, NasC2O.u, and sodium acetate, 
NaC2H;0O2, respectively, will be neutral toward phenol phthalein 
or litmus. The experiment shows that while the three acids 
differ very greatly in “ strength ”’ as shown by a reaction which 
depends on the number of hydrogen ions, H*, they give at a 
given dilution, they are nearly equal in their power of neutral- 
ization, which depends for these acids and the indicator chosen 
on the number of hydrogen ions which can be obtained from 
them by complete ionization. It should be remembered, how- 
ever, that sodium oxalate and sodium acetate are not, strictly 
speaking neutral salts. The reason for this lack of neutrality 
will be clearer after a study of the following paragraph. 

Use of Indicators. It has been stated (p. 122) that indicators 
are colored substances which exist in two forms, one of which 
is stable in the presence of hydrogen, H*, ions while the other 
is stable in the presence of hydroxide, OH~, ions. A neutral 
solution is one in which the numbers of hydrogen and hydroxide 
ions are equal. It has been pointed out in the table on p. 383 
that the number of mols of water ionized in one liter is 0.0000001, 
or 10°77 at 25°. Since in the ionization reaction 

TLOzatite OHe 
we must have 
Cut X Con- = constant 


so long as the solution consists mostly of water, it follows that 
in dilute solutions 


Cyt >< Con- = 1077 x 1077 = 10-4 


If an acid is added to water in sufficient amount to increase 
the concentration of the hydrogen ions from 1077 to 10°, the 
concentration of the hydroxide ions must fall to 1078. Only 
0.0000009 gram-mol of hydrogen ions would be required to 
produce such a change in a liter of water, and this would be 
given by about 0.001 cc. of tenth-normal hydrochloric acid, 
while 0.01 ce. would increase the hydrogen ion concentration 
bon) 6 


A TEXTBOOK OF CHEMISTRY 


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USE OF INDICATORS 389 


Indicators are weak acids or bases, and the change in color 
does not usually occur in an exactly neutral solution. The table 
on p. 388 gives the concentrations of hydrogen and hydroxide 
ions at which the change in color occurs for some of the more 
common indicators. 

It would seem, at first thought, that only an indicator which 
changes exactly at the neutral point would be suitable, but this 
is by no means true. In titrating a strong acid with a strong 
base a few hundredths of a cubic centimeter of tenth-normal 
acid or alkali will carry the concentration of hydrogen or hy- 
droxide ions so far to one side of the neutral point that any of 
the indicators for which the acidity is between 10~° and 10° 
will give a sharp end point. 

In titrating a weak acid, such as acetic acid, HC:H3Os., 
with a strong base, such as potassium hydroxide, KOH, as the 
neutral point is approached the potassium acetate, KC2H;Oz, 
formed is much more highly ionized than the acetic acid and 
the acetate ions carry the ionization : 


HC.H;0. ss He + C2H30.— 


far to the left. The number of hydrogen ions then becomes 
very small, while a considerable amount of acetic acid is still 
unneutralized. Under such conditions the change in color for 
methyl orange or methyl red may appear before the acid is 
completely neutralized and the end reaction will not be sharp, 
1.e. the change in color will appear gradually during the addition, 
sometimes, of a cubic centimeter or more of the alkali. But if 
phenol phthalein is used, the change in color will not occur till 
the true neutral point is passed and then a very slight excess 
of alkali will carry the concentration of the hydroxide ions far 
beyond the neutral point. The end will be sharp and will cor- 
respond closely to the exact neutralization of the acid. 

When a ‘weak base such as ammonium hydroxide, NH,OH, 
is titrated with a strong acid, as hydrochloric acid, HCl, the 
conditions are reversed and such indicators as methyl red, 


390 A TEXTBOOK OF CHEMISTRY 


methyl orange or cochineal, which change color in a faintly 
acid solution, are most suitable. 

With very weak acids or bases the hydrolysis of the salts 
formed may carry the acidity or alkalinity of the normal salt 
far to one side of the true neutral point. An illustration of the 
use of indicators in such a case will be given later (p. 464). 

Systematic Study of the Metals. In discussing the metals it 
is natural to consider first, as with the non-metals, their prepa-~ 
ration and properties. After this their compounds may be dis- 
cussed in the same order which has been followed in the consid- 
eration of the nonmetallic elements: oxides and hydroxides ; 
chlorides, hypochlorites, chlorates, bromides, iodides, fluorides ; 
sulfides, sulfites, sulfates; nitrites, nitrates, phosphates, arsen- 
ates; carbonates, bicarbonates, salts of organic acids, cyanides ; 
silicates; borates. Most of the metals form nearly all of these 
classes of compounds, but only those compounds which are of 
some particular scientific, historical or industrial importance can 
be mentioned in this book. 

Before taking up the individual metals it seems desirable to 
give some general statements with regard to the preparation or 
metallurgy of the metals and about the preparation and proper- 
ties of the various classes of compounds. 

Metallurgy. The first metals used by man were those which 
are found free in nature, such as copper, silver and gold. The 
use of these metals and their alloys marks the beginning of the 
“Bronze Age,’ which reaches back into prehistoric times — 
very late and recent when we consider the long history of the 
race, but early in a period when mankind became organized 
in society and differentiated sharply from the animal world. 
The discovery of methods of reducing iron from its ores by means 
of wind furnaces probably occurred during the period of history 
which has been recorded in inscriptions, but still so early that 
no definite record is to be found. Inscriptions in Egypt show 
that iron was made there at least 3000 years ago. * From that 
period till the middle of the nineteenth century all methods for 
industrial metallurgy depended on the use of charcoal, coal or 


METALLURGY O91 


other fuels containing carbon. The most important of these 
methods depended on the direct reduction of oxides of the metals 
by means of these fuels, while a few depended on the roasting 
of a sulfide (mercury, p. 485) or on the roasting of a sulfide 
followed by reduction through the interaction of an oxide with a 
sulfide or sulfate (copper, p. 427, and lead, p. 513). 

The first preparation of aluminium by Deville in 1854 in 
sufficient amount to demonstrate its valuable properties led to a 
strong desire to secure cheap sodium for use in its production. 
This resulted in the development of the reduction of sodium 
carbonate to metallic sodium by means of carbon and the use 
of the latter for the preparation of aluminium from its chloride, 
AICl;. But this method of preparing aluminium did not attain 
any considerable industrial. importance. 

Electrolytic methods were used for the deposition of thin 
films of copper and other metals as early as 1836, but electrical 
methods for preparing and refining metals could not be used 
on a large scale till the development of dynamos during the 
last quarter of the nineteenth century made the production of 
relatively cheap electrical energy possible. 

The first industrial use of an electric furnace seems to have 
been its application to the manufacture of aluminium bronze by 
the Cowles Brothers of Cleveland, Ohio, in 1884. Their labora- 
tory experiments with the method began in 1882. The Hall 
method for the electrolysis of aluminium oxide dissolved in 
eryolite soon rendered the Cowles furnace industrially worth- 
less for this particular purpose, but electric furnaces are now 
extensively used for metallurgical processes and for many other 
forms of chemical manufacture. 

The relatively cheap electrolytic manufacture of aluminium 
has not only given an abundant supply for use in the metallic 
form, but has led to the development of processes for the pro- 
duction of chromium and other metals by heating their oxides 
with aluminium. 

It will be seen from this brief sketch of the historical develop- 
ment of metallurgy that many important processes are of very 


392 A TEXTBOOK OF CHEMISTRY 


recent origin. It is only a very few years since several metals 
which now have very important industrial uses, were scarcely 
more than scientific curiosities. Further rapid development 
along such lines is to be confidently expected. 

Oxides. All metals, without exception, may be combined 
with oxygen and nearly all metals combine with the element 
directly at suitable temperatures. Even the so-called ‘ noble ” 
metals, gold, silver and platinum combine with oxygen under 
some conditions, but the oxides are very unstable, are easily 
decomposed by heat alone and are reduced by hydrogen at 
ordinary temperatures. 

A few metals, especially sodium, potassium and barium, com- 
bine with oxygen to form peroxides, in which two oxygen atoms 


O 
are united, as in Bat |, but in nearly all of the metallic oxides 


O 
the oxygen seems to be united only with the metal, and the 
valence of the metal is apparent from the formula of the oxide. 
Nearly all nitrates, carbonates and hydroxides are decom- 
posed by heat with the formation of oxides: - 


CaCO; — CaO -+- CO,z 
Cu(OH),> = CuO + H.O 


It is doubtful if any oxide dissolves appreciably in water as 
an oxide, and those oxides which do not combine with water 
to form hydroxides are practically insoluble. 

Hydroxides. ‘The alkali metals, sodium, potassium, etc., 
and the alkali earth metals, calcium, strontium and barium 
decompose water with the formation of hydroxides at ordinary 
temperatures : 


Na + HOH’ = NaOH + H 
Ca + 2HOH = Ca(OH). + 2H 
Magnesium decomposes water at 100°. Other metals, as zine 


and iron, which decompose water at higher temperatures, form 
oxides instead of hydroxides, though it is, of course, possible 





SOLUBILITY OF SALTS 393 


that the latter are formed at first and immediately decom- 
posed. 

Iron is converted by the combined action of water and air into 
iron rust, a combined oxide and hydroxide having the composi- 
tion of the mineral limonite, Fe,03.2Fe(OH)3. 

Practically all hydroxides of the metals except those of the 
alkali and alkali-earth metals are insoluble in water. For this 
reason the hydroxides of nearly all other metals are precipitated 
from solutions of their salts by solutions of sodium or potassium 
hydroxide. In a few cases, especially those of silver, cuprous 
copper, mercurous and mercuric mercury, these hydroxides 
precipitate an oxide instead of the hydroxide, doubtless because 
the hydroxides of these metals are unstable: 


AgNO; + NaOH = [AgOH] + NaNO; 
[2AgOH] =Ag.0 + H.0 


Solubility of Salts. Practically all salts of the alkali metals 
(lithium, sodium, potassium, ammonium) are soluble in water, 
the only important exceptions being sodium pyroantimonate, 
NazH2Sb20;.6 H2O, potassium and ammonium chloroplatinates, 
K.PtCl, and (NH4)oPtCls,! potassium perchlorate, KClO,, and 
potassium cobaltinitrite, K3;Co(NOze)s, or potassium = silver 
cobaltinitrite, K,AgCo(NQOz)s. Some of these are, however, 
more soluble than those salts of other metals which are usually 
counted as insoluble. There is probably no salt which is 
wholly insoluble in water. 

Nearly all salts of the strong monobasic and bibasic acids 
are also soluble. This includes chlorides, bromides and iodides, 
fluorides, chlorates and perchlorates, sulfites and sulfates, ni- 
trites and nitrates. The most important exceptions are the 
chlorides, bromides and iodides of .silver, cuprous copper, mer- 
curous mercury and lead, AgCl, AgBr, AgI, CueCle, Cuele, 
HgeCle, Hgele, PbCle (slightly soluble), PbBre, and PbIs, mer- 


1 Rubidium and cesium chloroplatinates, RbePtCls and Cs:PtCl. 
are still less soluble. 


394. A TEXTBOOK OF CHEMISTRY 


curic iodide, Hgle, calcium fluoride, CaF:, barium sulfite, BaSOs, 
and calcium, strontium, barium and radium sulfates, CaSO, 
(slightly soluble), SrSO1, BaSO., RaSOx. 

Normal salts of phosphoric, H3POx, arsenious, H3AsQOs, arsenic, 
H3AsOu, carbonic, H2CQs, silicic, H2SiOs, etc., and boric, H3BOs, 
acids, with the exception of those of the alkalies, are insoluble. 
Sulfides other than those of the alkalies are either insoluble in 
water or are hydrolyzed with the formation of a hydrosulfide 
(as Ca(SH).2), and a hydroxide, or of hydrogen sulfide, H.S, and 
an insoluble hydroxide, such as Al(OH)s. 


CHAPTER XXIII 


ALKALI METALS: LITHIUM, SODIUM 


General Properties of the Alkali Metals. The alkali metals 
are univalent elements which combine with hydroxyl to form 
the strongest bases, hydroxides which are easily soluble in water 
and which are largely ionized in solutions of moderate concen- 
trations. ‘They are the most active of the metallic elements, 
decomposing water rapidly at ordinary temperatures and tar- 
nishing almost instantly in ordinary air, owing to the formation 
of a film of hydroxide. 

Their affinity for the halogens is also so strong that sodium 
and potassium have often been used to decompose halides for 
the preparation of metals and other elements. 

As has been stated in the last chapter, nearly all salts of the 
alkali metals are soluble in water. Normal salts of weak acids, 
such as the sulfides, carbonates, cyanides, phosphates, silicates 
and borates are hydrolyzed by water, and their solutions have a 
strongly alkaline reaction. 

The metals of the group have a low specific gravity, lithium, 
sodium and potassium being lighter than water. Their melting 
points range from 186° for lithium to 26.5° for cesium and the 
boiling points, from above a red heat for lithium and 742° for 
sodium to 270° for cesium. 

Lithium, Li, 6.94, is usually considered as one of the rarer 
elements and of comparatively little importance. It is found 
in a number of silicates and in small amount in practically all 
natural waters. The metal has a specific gravity of only 0.51. 
Hydrogen and helium are the only elements which are lighter 
than lithium when in the solid or liquid state. The metal com- 
bines with hydrogen to form the hydride, LiH, and with nitrogen 

395 


396 A TEXTBOOK OF CHEMISTRY 


to form the nitride, LisN. It has been used in the separation of 
argon from the atmosphere because of its strong affinity for 
nitrogen. 

Lithium carbonate, LigCO3, may be decomposed to lithium 
oxide, LigO, and carbon dioxide by heating it in a current of 
hydrogen. Both the carbonate and the phosphate, LisPOx,, are 
difficultly soluble in water. In these properties lithium re- 
sembles magnesium, the second element of the second group, 
rather than the other alkali metals. Beryllium, the first ele- 
ment of the second group, approaches aluminium in its proper- 
ties In a similar manner. 

Lithium Urate, LiC;H303N,, is soluble in water, and this fact 
led physicians to the belief that the administration of lithium 
carbonate or of natural waters containing lithium would be 
beneficial to patients suffering from rheumatism or gout and 
they have been much employed as remedies in those diseases. 
A more careful study has shown that these compounds are worth- 
less for such a purpose, but the ingestion of large quantities of 
water with the lithium compounds probably exerts a beneficial 
effect. | 

Lithium compounds impart to the Bunsen flame a brilliant 
red color and give a spectrum of two red lines, one of which is 
very bright. 

Atomic Weight of Lithium. Law of Dulong and Petit. It 
has been pointed out (p. 92) that the most satisfactory method 
of selecting the true atomic weight of an element consists in 
finding the weight of the element contained in a gram-molecular 
volume (22.4 liters at 0° and 760 mm.) of that gaseous compound 
which contains the smallest quantity of the element in this 
volume. But lithium forms no compound whose weight in the 
gaseous form has been determined, and a considerable number 
of other elements form no compounds which can be converted 
into gases without decomposition. The atomic weights of such 
elements must, of course, be selected in a different manner. 
For this purpose the law of Dulong and Petit, discovered in 
1819, has been useful. ‘These chemists found that the quantity 


LAW OF DULONG AND PETIT 397 


of heat required to raise the temperature of one gram-atom of 
an element one degree is approximately 6.6 calories. If this 
quantity of heat is applied to 7 grams of lithium or to 65 grams 
of zinc or to 200 grams of mercury, it will, in each case, raise the 
temperature one degree. 

The law is also frequently stated that the specific heat of an 
element multiplied by its atomic weight is a constant quantity. 
The following table will make this clear: 











ELEMENT ae ore Sp. Ht. X At. Wr. 
ritnium ) =. ARE Ss 0.94 fé 6.6 
Graphite (at 11°) Re Ad ae ys 0.16 ee 1.9 
rapuitenat Gide y) co, |e 0.467 12; 5.6 
DInCONe eC Serkee sy 52) tra 0.16 28.4 4.5 
alCHuINeey ee nel. Sh a vol. Oal7, 40. 6.8 
‘Sas he op es gee Pere 0.093 65.4 6.1 
DPOMINC Opa Pars "aLe ey, tae ee 0.084 80. 6.7 
VLGTCULY kil. Seals rr 5 0.033 200. 6.7 
CAC Mh ieee? We Neg s | Brake pt: 0.03 207: 6.2 











It will be seen from the table that graphite and silicon depart 
rather widely from the law, though the former approaches it 
‘more closely at high temperatures. All of the metallic elements 
and all elements having atomic weights above 40 conform ap- 
proximately to the law. The law is at best, however, only 
approximate and is of service only in selecting between rather 
widely divergent possible values for an atomic weight. Thus 
the atomic weight of calcium might be 20, 40 or 60, according 
as the formula of the chloride is CaCl, CaCly or CaCl3. But of 
these three values only an atomic weight of 40 agrees with the 
law. 

The laws of Avogadro and of Dulong and Petit have usually 
been considered as independent and wholly unrelated.! A 


1 See, however, G. N. Lewis, J. Am. Chem. Soe. 29, 1165 and 
1516 (1907). 


398 A TEXTBOOK OF CHEMISTRY 


little consideration, however, shows us that if we accept the 
kinetic-molecular theory, this is not the case. At foundation 
Avogadro’s law depends on the fact that molecules of different 
weights exchange energies, when in collision with each other 
or with the walls of the containing vessel at a given tempera- 
ture, in such a manner that the average value of } mv (m = 
mass, 0 = velocity) is constant and is independent of the 
weight of the molecule. The law of Dulong and Petit must 
depend on a similar property of the atoms of the elements in 
the solid or liquid state. 

The Quantum Theory. A theory of molecular, atomic and 
radiant energy, called the quantum theory, has been developed 
by Plank, Einstein, Nernst, Sackur, Debye, Sommerfeld, Bohr 
and others. 

The theory supposes that radiant energy is always given out 
or absorbed in definite units, called quanta. It has been most 
successful in the explanation of the spectral lines of hydrogen, 
helium and other elements. It is supposed that an electron 
rotating about the positive nucleus of an atom may shift from 
one orbit to another. If it falls from an outer orbit to one 
closer to the nucleus it gives out one or more quanta of radiant 
energy. ‘These are transmitted through space and give the char- 
acteristic spectral lines of the element. If radiant energy of 
the same period passes an atom an electron may be lifted to an 
orbit further from the nucleus and the energy will be absorbed. 
This gives the dark lines of the solar spectrum. 

Sodium, Na, 23. Sodium chloride, NaCl, or common salt, 
forms about 75 per cent of the residue left by the evaporation 
of sea water. It is also found in enormous beds of rock salt in 
Germany, Louisiana, Kansas, Utah and elsewhere and in 
strong brines found by boring deep wells in very many places. 
Sodium is taken up by seaweeds in their growth very much as 
potassium is taken up by land plants, and the ashes of sea- 
weeds contain considerable quantities of sodium carbonate. 
Deposits of sodium sesquicarbonate, NaHCO3.Na2CO3.2 H2O, 
called trona, of sufficient extent to be of industrial importance 
are found in Egypt and in Venezuela. The occurrence of 


SODIUM _ 399 


borax, Na2B,407.10 H2O, in lakes in California has been men- 
tioned; also that of sodium nitrate, NaNO3. These are of 
value, primarily, for the boron and nitrogen which they contain. 
Sodium is a constituent of practically all silicious rocks. 

Metallurgy. Properties. Metallic sodium and_ potassium 
were first prepared by Sir Humphry Davy in London in 1807 
by the electrolysis of moist sodium and potassium hydroxides. 
The discovery awakened very great interest, both because the 
metals showed very striking properties, quite different from 
those of any metals hitherto known, and because it indicated 
very clearly that many other earthy substances contain ele- 
ments which could not at that time be prepared in the free 
state. Within a few years the new metals proved effective 
agents for the liberation of a number of other elements. 

Sodium may also be prepared by heating a mixture of sodium 
carbonate and carbon: 


NaeCO3 + 2C = 3CO + 2Na 


The sodium, which boils at 742°, distills from the mixture and 
is collected in iron condensers. During comparatively recent 
years metallic sodium is prepared commercially by various 
electrolytic methods, from the hydroxide, the nitrate or the 
chloride. 

Sodium is a silver white metal which tarnishes instantly on 
exposure to moist air. In dry air at 300° to 400° it is oxidized 
to sodium peroxide, NagO». It is kept in sealed cans or under 
kerosene to protect it from the action of the air. It melts at 
97.5° and boils at 742°, giving a dark green vapor. The specific 
gravity of the solid is 0.97. When a small piece of sodium is 
thrown on water the heat of the reaction causes it to melt and 
the globule of metal rolls rapidly over the surface of the water 
without taking fire or igniting the hydrogen, differing in this 
respect from potassium. If the metal is thrown on a piece of 
filter paper floating on the water, the heat is concentrated and 
the hydrogen takes fire and burns with a yellow flame. In 
both cases, of course, sodium hydroxide is formed. 


400 A TEXTBOOK OF CHEMISTRY 


Sodium is manufactured in considerable amounts for use in 
preparing sodium peroxide, Na2Os, for the preparation of a 
mixture of potassium and sodium cyanides from potassium ferro- 
cyanide (p. 319), and for use in the synthesis of a variety of 
organic compounds. 

The Alkali Industry. Sodium is an essential constituent of 
common soap, of glass, of salsoda, or washing soda, and of 
baking soda. As common salt, NaCl, is much cheaper than 
any other compound of sodium, it now furnishes the basis for 
the preparation of all of these substances, but it is necessary to 
prepare from it, at first, one of the sodium carbonates or sodium 
hydroxide. ‘Till the close of the eighteenth century an impure 
sodium carbonate from Egypt and the ash of seaweeds were 
used as the sources of sodium carbonate and sodium hydroxide 
for the manufacture of hard soap, while potassium carbonate 
from wood ashes was also extensively used for the manufacture — 
of soft soap. During the disturbed commercial relations which 
followed the French Revolution, the foreign supplies of sodium 
carbonate were cut off and all available potassium compounds 
were needed for the manufacture of gunpowder. This caused 
the French government to offer a prize for a method of manu- 
facturing sodium carbonate from salt. The prize was awarded 
to Leblanc and his process was used for a short time in France, 
but could not there compete with the sodium carbonate from 
other sources when commercial relations with other countries 
were again established. Leblanc himself did not secure any 
permanent advantage from his invention and died in a poor- 
house. About twenty years later Musgrave, in England, took 
up the process again and succeeded in making it a commercial 
success. It held the field of alkali manufacture almost without 
competition for fifty years. About 1860 it had to meet the 
competition of the ammonia-soda process, the principles of 
which had been discovered in 1838, but which was first put 
into successful operation by Solvay more than twenty years 
later. From then till the close of the nineteenth century the 
Leblanc process continually lost ground in competition with the 


SODIUM HYDROXIDE eAGl 


Solvay manufacture, maintaining a precarious existence only by 
the most careful conservation of the by-products, hydrochloric 
acid or chlorine and sulfur. In 1900 only two large Leblanc fac- 
tories remained in the world, one in England and one in Germany. 

The most serious difficulty with the ammonia-soda process is 
that the chlorine of the salt is converted into calcium chloride 
or magnesium chloride, practically worthless products. In the 
closing years of the nineteenth century, with the stimulus of 
comparatively cheap electrical energy, many electrolytic pro- 
cesses were developed which give both chlorine and sodium 
hydroxide directly from salt. It seems probable that these 
processes will eventually displace the Solvay process, at least 
for the production of caustic alkali. 

Sodium Hydroxide. So long as trona from Egypt or the ashes 
of sea plants were used, sodium hydroxide was prepared by 
treating a solution of these ,with slaked lime. A more pure 
sodium hydroxide was prepared in the same way from the 
sodium carbonate of the Leblane or ammonia-soda processes : 


Na2CO; + Ca(OH). = 2 NaOH + CaCO; 


The reaction depends, of course, on the insolubility of the 
calcium carbonate. In very concentrated solutions the reaction 
may be reversed because calcium hydroxide, Ca(OH)s, is also 
difficultly soluble, and with a high concentration of hydroxide 
ions, OH’, the solubility product for that substance may be 
exceeded even in a solution containing so insoluble a salt as 
calcium carbonate. 

As has been stated above, sodium hydroxide is now prepared 
on a large scale by electrolysis. Many different patents have 
been issued for such processes and it is probably too soon to 
decide which forms are likely to prove permanently suitable. 
In some forms a diaphragm, usually of asbestos, is used to sepa- 
rate the anode space from the cathode. The anode must be of 
carbon or of platinum or platinum-iridium and the anode space 
is inclosed so that the chlorine liberated may be collected and 
utilized. The cathode is usually of iron. Chlorine is liberated 


402 A TEXTBOOK OF CHEMISTRY 


at the anode, while hydrogen is liberated at the cathode and 
sodium, Nat, and hydroxide, OH~, ions remain in solution, 
the hydrogen, of course, coming from the water, though the 
transfer of ions through the solution is mainly that of sodium 
and chloride ions. The points aimed at are to secure as high a 
concentration of hydroxide, OH, and as low a concentration 
of chloride, Cl-, ions as possible in the cathode space and the 
reverse of this around the anode. ‘To this end the salt solution 
is continuously introduced at the anode while the hydroxide 
solution is removed from the cathode. The presence of hy- 
droxide at the anode leads to the formation of hypochlorite and 
loss of current. ‘The hydroxide solution may be concentrated 
till nearly all of the salt, NaCl, remaining in it separates, leav- 
ing a solution in which nearly all of the sodium is in the form of 
hydroxide. This solution is then evaporated till the water has 
been expelled, which requires a comparatively high temperature. 
The sodium hydroxide obtained in this way is sufficiently pure 
for the manufacture of soap and for many other industrial uses. 

The Castner-Kellner apparatus gives an almost pure solution 
of sodium hydroxide directly. It consists of a slate box divided 
into three com- 
partments by two 
partitions, which 
fit only loosely in 
grooves in the 
bottom of the box 
(Fig. 94). Mer- 
cury placed on the bottom of the box seals these, giving a contin- 
uous metallic layer for the three compartments, but prevents a 
dilute solution of sodium hydroxide placed in the central com- 
partment from mixing with the brine placed in the two side 
compartments. Graphite anodes are placed in the two side 
compartments and an iron cathode in the central one. Chlo- 
rine is evolved from the anodes and is, of course, collected and 
used for the manufacture of bleaching powder or for some 
other purpose. The mercury in the two side compartments is 





SODIUM HYDROXIDE 403 


negative as compared with the graphite anodes and the sodium 
liberated at its surface combines with it to form a liquid sodium 
amalgam. By a slight tilting motion the amalgam is caused to 
flow alternately to one side or the other and so is brought 
into the central compartment. Here it is positive with refer- 
ence to the more negative cathode and the hydroxide ions 
brought to its surface by the current combine with the sodium 
of the amalgam to form sodium hydroxide, while the hydrogen 
ions of the water are discharged and liberated as free hydrogen, 
Hy, at the surface of the iron cathode. The hydroxide solu- 
tion is kept at a constant concentration by introducing water 
at one side and removing some of the solution at the other. 
Salt is added from time to time to the side compartments. 

Sodium hydroxide is a white solid, which melts at a red heat. 
It deliquesces on exposure to the air, but the solution soon ab- 
sorbs carbon dioxide and then evaporates, leaving a residue of 
sodium carbonate. 

Sodium hydroxide is used in the manufacture of ordinary 
hard soaps and may be used for the preparation of many of 
the salts of sodium. 

The specific gravity of solutions of different concentrations 
is as follows: 


DeEnsITY OF SOLUTIONS OF SopIuM HYDROXIDE 

















PER CENT GRAMS OF NaOH 
SPECIFIC GRAVITY NaOH rw 1000CG, 
1.0555 5) O.cd 
LA00t 10 Diag 
1.1665 15 17.497 
1.2219 20 24.438 
122771 25 31.928 
1.3312 30 39.936 
1.3838 30 48.433 
1.4343 40 57.372 
1.4828 45 66.726 


1.5308 50 76.515 











404 A TEXTBOOK OF CHEMISTRY 


Sodium hydroxide dissolves in water with the evolution of 
considerable heat, and it will be seen from the table that the 
addition of a small amount of sodium hydroxide causes the 
volume of the water to diminish. 

Sodium Oxide, Na.O. With the exception of lithium oxide, 
Li,O, the oxides of the alkali metals cannot. be prepared by 
heating the hydroxides or carbonates. In this respect they 
differ from all other metallic oxides. Sodium oxide may be 
prepared by heating sodium hydroxide, NaOH, with metallic 
sodium. It combines with water to form the hydroxide, but 
has, at present, no practical importance. 

Sodium Peroxide, Na.O.; is prepared by heating metallic 
sodium to 300°—400° in dry air. The sodium is placed in shallow 
aluminium trays, which are passed slowly through long ovens 
one way, while air passes in the opposite direction. In this 
way the pure sodium comes at first in contact with air which 
has been deprived of most of its oxygen and a too vigorous 
action is avoided, while the action is finally completed at the 
other end with fresh air. 

The use of fused sodium peroxide containing a little copper 
oxide, under the name of “ oxone,” for the preparation of 
oxygen has been mentioned (p. 21). It may be converted by 
cold, moist air into a hydrate, Na2O2.H2O, which can be dis- 
solved in water with little decomposition. It is hydrolyzed, how- 
ever, to sodium hydroxide, NaOH, and hydrogen peroxide, H2Os. 

On treatment with cold, dilute acids sodium peroxide gives a 
solution of hydrogen peroxide, H2O2, which is used to bleach 
silk, wool, hair and other substances which would be affected 
injuriously by chlorine. 

Sodium peroxide is also a very valuable oxidizing agent for 
many laboratory uses. 

Sodium Chloride. Salt is sometimes obtained by direct 
mining, but rock salt is rarely sufficiently pure for direct use, and 
it offers especial difficulties in mining, owing to its effect in 
dulling tools used to cut it and because blasting does not loosen 
it up satisfactorily. It is found better to prepare an opening 


SODIUM CHLORIDE 405 


in the bed of salt and allow water to stand in contact with it 
till a saturated solution is obtained, many of the impurities 
present remaining undissolved and settling to the bottom. 
The solution is then pumped out and evaporated to crystallize 
the salt. As salt is nearly as soluble in cold as in hot water, it 
cannot be crystallized practically by cooling a hot solution. 
In some places, especially at Syracuse, New York, and in Michi- 
gan, saturated brines are obtained directly from artesian wells. 

For the evaporation of the brines triple-effect evaporators 
are used to advantage. The principle of these is shown in the 
accompanying diagram (Fig. 95). The brine in pan A is heated 
directly, or by superheated steam in coils or in a false bottom, 





Fig. 95 


and it boils under atmospheric pressure. The steam from this 
pan passes under B, in which a pressure of perhaps 550 mm. is 
maintained so that the condensation of the steam from the first 
pan beneath it will cause the brine which it contains to boil. 
The steam from this will, in turn, cause the brine in C to boil 
under a pressure of 300 mm. By such an arrangement a given 
weight of coal will evaporate nearly three times as much water 
as it would if used directly in the usual manner.! 

Commercial salt contains small quantities of various impuri- 
ties. The most objectionable, perhaps, is magnesium chloride, 

1Tt may be remarked, incidentally, that when this process is 
used for the preparation of distilled water, from waters containing 
relatively small amounts of solids in solution and the successive 
differences of pressure may be much less, as many as ten boilers 
may be used in series. For a description of the Yaryan Evapora- 


tor, which uses a modification of this system, see J. Soc. Chem. 
Ind. 14, 112 (1895). 


406 A TEXTBOOK OF CHEMISTRY 


which makes it hygroscopic or even deliquescent in moist air. 
Pure sodium chloride can be obtained by precipitating a solu- 
tion of salt with concentrated hydrochloric acid. 

Sodium chloride crystallizes in cubes. It melts at 820° and 
may be volatilized at a high temperature. The crystals usually 
decrepitate on heating, owing to water inclosed in them. 

Salt is an essential constituent of human diet, furnishing 
chlorine for the hydrochloric acid of the gastric juice. 

Sodium Sulfate. Glauber’s Salt, NagSO4.10 HO. By heat- 
ing salt with the theoretical amount of sulfuric acid it may be 
converted almost quantitatively into anhydrous sodium sulfate, 
Na2SO,. The operation is carried out on a large scale as the 
first step in the Leblanc soda process. The anhydrous sulfate 
is also used in the manufacture of glass. 

The crystallized hydrate, Na2SO.1.10 H.O, has a solubility 
in water which increases very rapidly with rising temperature 
till the transition point, 32.383°, is reached. If crystals of the 
hydrate are heated above this temperature, they are transformed 
into a mixture of anhydrous sodium sulfate and a saturated 
solution of the latter, which is less soluble than the crystallized 
hydrate at temperatures above this point. The transition is 
accompanied by an absorption of heat in very much the same 
manner as the melting of ice and may be used as an accurate, 
fixed point for the correction of thermometers. (Richards, Z 
physik. Chem. 43, 465.) 

* The transition point is a quadruple “i in the nomenclature 
of the phase rule (p. 107), the four phases being water vapor at 
a pressure of 30.8 mm., the hydrate, Na2SO.1.10H2O, the 
anhydrous salt, Na2SOu., and the saturated solution. As there 
are four phases and only two components, sodium sulfate and 
water, the system is invariant and there can be no change in 
temperature or pressure without the disappearance of one of the 
phases. 


1 Practically, the transition point is determined in contact with 
air at atmospheric pressure and the effect of pressure is not con- 
sidered. In the presence of the vapor phase only, the temperature 


SODIUM SULFATE 407 


The solubility of sodium sulfate is shown in the diagram, Fig. 
96. The concentrations are given for the anhydrous salt through- 
out. Below the transition point the hydrate separates on evap- 
orating or cooling the solution, though supersaturated solutions 


Solubility—Grams of Na,SO,in 100 grams of water 





soinjerodue J, 





of the transition point would be slightly different, just as the true 
transition point for water, ice, vapor is 0.0076° higher than the melt- 
ing point of ice under atmospheric pressure, which is used. as the 
zero point for thermometers. 


408 A TEXTBOOK OF CHEMISTRY 


are easily formed (p. 80). Above the transition point the anhy- 
drous salt separates on heating or on evaporation. 

Sodium sulfate is one of the chief active constituents of 
Hunyadi water and of some other similar medicinal waters. 

Acid Sodium Sulfate or Sodium Bisulfate, NaHSQ,, is formed 
as the first step in the preparation of the sulfate from salt, or 
in the manufacture of nitric acid from sodium nitrate, NaNO3. 
At about 300° it loses water and is converted into sodium pyro- 
sulfate, NaoS2O7. At a still higher temperature it is decom- 
posed into sodium sulfate and sulfur trioxide, SO3. Sodium 
pyrosulfate is used in analytical chemistry to dissolve aluminium 
oxide, Al.Os, ferric oxide, Fe.O3, and titanium oxide, TiO», after 
these have been brought to a difficultly soluble form by ignition. 
(Hillebrand, Analysis of Silicate and Carbonate Rocks, Bulletin 
422, U.S. Geol. Survey, p. 105.) 

Sodium Sulfite, NasSO3.H:,O, is prepared by burning sulfur 
and passing the sulfur dioxide formed through a solution of 
sodium carbonate. It is used as a reducing agent in photog- 
raphy and for the preparation of sodium thiosulfate. 

Acid Sodium Sulfite, or Sodium Bisulfite, NaHSOs, is pre- 
pared by passing sulfur dioxide in excess into a solution of sodium 
carbonate. It is sometimes used as an addition to cider to stop 
fermentation. A 40 per cent solution is very conveniently 
used for the preparation of sulfur dioxide in the laboratory. 

Sodium Hyposulfite, Na.S.O,. A solution of this salt is 
prepared by the action of zinc on a solution of sodium bisulfite, 
NaHSOs, containing an excess of sulfurous acid, H2SO:. 


2 NaHSO; + H.SO3 aL Zn = NaeS.O, -}- ZnSOz + yd H.O 


Sodium hyposulfite is a powerful reducing agent and is used 
especially to reduce indigo to indigo white (p. 341). 

Sodium Thiosulfate, Na:S.O3.5 H.O, is prepared by dissolv- 
ing sulfur in a solution of sodium sulfite. It dissolves the halides 
of silver (AgCl, AgBr and AgI) and is used to fix photographic 
pictures (p. 445). It is usually called by photographers and 
pharmacists “sodium hyposulfite.” It has also been used in 


SODIUM SULFIDE 409 


extraction of silver from its ores in the so-called ‘ hyposulfite- 
lixiviation ”’ processes. 

The anhydrous thiosulfate, Na2S203, can be obtained by dry- 
ing the crystals at a moderate temperature. If the dry salt is 
heated, it decomposes into sodium sulfate, sodium sulfide, and 
sulfur or sodium polysulfide. 

If a solution of sodium thiosulfate is warmed with copper 
sulfate, cuprous sulfide, Cu2S, and sulfur are precipitated, while 
sodium sulfate remains in solution. These reactions show that 
the compound retains the reducing properties of the sulfites 
and also the properties, in part, of a sulfide. This recalls the 
method of preparation and agrees well with the formula 


Na—S Ze O 
XK , which is assigned to the compound. 
Na O 


Sodium Tetrathionate, Na2S,0..H20, is formed by the action 
of iodine on a solution of sodium thiosulfate : 


me NapS.O3 +. I, = NaeS4Og + 2, Nal 


The reaction is quantitative and is much used in volumetric 
analysis. 

Sodium Sulfide, Na.S, may be prepared by passing hydrogen 
sulfide into a solution of sodium hydroxide in the requisite 
amount and evaporating the solution to dryness with exclusion 
of air. It is hydrolyzed by water, giving a strongly alkaline 
solution. 

If a mixtdre of sodium carbonate and sulfur is heated, or if 
any metallic sulfide is heated with sodium carbonate on charcoal 
or if any metallic sulfate is heated with sodium carbonate in the 
reducing flame (p. 304)-on charcoal, a sulfide is formed. If this 
is moistened with water on a silver coin a black spot of silver 
sulfide, AgeS, will be formed. The reaction is used as a test 
for sulfur in any form of inorganic combination. 

Sodium Hydrosulfide, NaSH, is formed when hydrogen sulfide 
is passed into a solution of sodium hydroxide in twice the amount 
necessary to form the sulfide. It loses hydrogen sulfide on evap- 


410 A TEXTBOOK OF CHEMISTRY 


oration of the solution with exclusion of air, in the same way that 
sodium bicarbonate, NaHCOQOs, loses carbon dioxide. 

Sodium Nitrate, NaNOs, is found in immense beds in Chile, 
South America. It has been the chief source from which nitric 
acid and saltpeter, KNO3;, have been prepared and large quan- 
tities have also been used for fertilizers, to furnish the nitrogen 
necessary for the growth of crops. It is estimated that the 
supply from Chile will be exhausted in a comparatively few 
years, but there seems now a good probability that the manu- 
facture of nitrates from the nitrogen of the air will soon be in a 
position to supply its place. 

Sodium nitrate crystallizes without water of crystallization 
in rhombohedra. It melts at 316°. It is hygroscopic and for 
that reason cannot be used in place of potassium nitrate for the 
manufacture of ordinary gunpowder, though its low molecular 
weight makes it, otherwise, more suitable. 

Sodium Nitrite, NaNOz, is prepared by heating sodium nitrate 
with metallic lead. It is very easily soluble, but crystallizes 
well. It is used in laboratories and in factories for the prepara- 
tion of diazonium compounds for the manufacture of dyestuffs 
and other important compounds. 

Sodamide, NaNH,, is prepared by passing ammonia over me- 
tallic sodium at 300°-350°. (See Dennis and Browne, J. Am. 
Ch. Soe. 26, 587.) It dissolves in liquid ammonia, ionizing to 
Na* and NH,° (p. 207). It is hydrolyzed by water to sodium 
hydroxide and ammonia. It has recently become important 
for the preparation of indigo. 

* Sodium Trinitride, NaN3, is formed by the action of nitrous 
oxide, N2O, on sodamide (p. 223). 

Disodium Phosphate, Nas.HPO,.12 HO, is the best known 
and most important of the phosphates of sodium. It is isomor- 
phous with the corresponding arsenate, Na:HAsO,.12 H;0, 
and usually contains some of that salt derived from the impure 
phosphorus or from the sulfuric acid used in the preparation of 
phosphoric acid. The salt is sometimes used in medicine as 
a mild cathartie. 


SODIUM CARBONATE ~ Sail 


Sodium Carbonate or Sal soda (Washing Soda), 
NayCO3.10 H:;O. The Leblanc Soda Process. The Leblanc 
soda process is carried out in three operations. 

1. Salt is treated with sulfuric acid on the hearth of a furnace 
which can be heated to complete the reaction and expel all of 
the hydrochloric acid. The latter is conveyed through a tower 
filled with coke over which water is trickling and the aqueous 
hydrochloric acid obtained is sold or used for the preparation of 
chlorine. In the early years of manufacture the acid was al- 
lowed to escape and produced disastrous effects upon vegetation 
in the neighborhood of the works. This led to stringent legis- 
lation forbidding the escape of the acid. Later, the recovery 
of the hydrochloric acid proved profitable, and this has been 
an important factor in preventing the complete abandonment 
of the process : 


2 NaCl + H2SO4 = Na.sO. + 2 HCl 


2. The sodium sulfate is mixed with charcoal, or coal, and 
limestone, CaCOs3, and heated to fusion in the “ black ash fur- 
nace.” ‘The mass melts, the sodium sulfate is reduced to sodium 
sulfide, NaoS, and the latter reacts with the calcium carbonate 
to form calcium sulfide, CaS, and sodium carbonate: 


Na.SO, -|- 2, C = NaoS -f- Ds CO,z 
Nass + CaCQ3 —— CaS + NasCO3 


3. The “ black ash,’ when cold, is leached with water, which 
dissolves the sodium carbonate and leaves most of the calcium 
sulfide undissolved. There is a tendency for the calcium sulfide 
to hydrolyze to calcium hydrosulfide, Ca(SH)2, and calcium 
hydroxide, Ca(OH)», but this takes place slowly and is repressed 
in the strongly alkaline solution, so that a fairly complete separa- 
tion is obtained. The sodium carbonate is obtained from the 
solution by evaporation and crystallization. If the evaporation 
and crystallization take place above 35.2°, the monohydrate, 
NasCO3.H.0, crystallizes from the solution. If the salt is crys- 
tallized below that temperature, by evaporation or by cooling 


412 A TEXTBOOK OF CHEMISTRY 


the solution, the dekahydrate, Na2CO3.10 HO, is formed. 
The temperature given, 35.2°, is, of course, the transition point 
from the dekahydrate to the monohydrate. 

The Leblane process, which was so important during the 
nineteenth century, has been completely displaced by other 
processes. 

Crystallized sodium carbonate is known as sal soda, or washing 
soda, and is used for laundry purposes, for the manufacture of 
soap and for softening water. Anhydrous, or calcined, sodium 
carbonate is used in the manufacture of glass and for the pre- 
paration of other compounds of sodium. 

Sodium carbonate is hydrolyzed by water to sodium bicar- 
bonate and sodium hydroxide, owing to the very trifling ioniza- 
tion of the hydrocarbonate ion, HCO37 : 


Nat + Nat + CO;- + Ht + OH- 
= Nat + Nat + HCO;- + OH- 


Sodium Bicarbonate, or Baking Soda, NaHCO;3. The Ammo- 
nia-soda Process. About 1860 Solvay succeeded in putting this 
process, which had been discovered many years before, into a 
successful form. It has been found more economical than the 
Leblanc process and by the close of the nineteenth century it had 
very largely displaced that method of manufacture. 

A strong brine is first treated with ammonia till it contains 
one molecule for each molecule of salt. A tower about 20 meters 
high and having a series of shelves is filled with the brine and 
carbon dioxide is forced into the tower at the bottom. Am- 
monium bicarbonate, NHizHCOs, is at first formed, but as sodium 
bicarbonate, NaHCQs, is the least soluble of the various com- 
binations of ions possible, it separates and is deposited in layers 
on the shelves: 


NHs3 ob H.O + CO, — NH,HCOs3 
NH.HCO; + NaCl = NH,.Cl + NaHCO; 


The sodium bicarbonate, after removal from the shelves and 
washing in centrifugals with a little water, is almost chemically 


SODIUM SILICATE 413 


pure. Sodium carbonate can be easily obtained from this by 
calcining it at a comparatively low temperature: 


2 NaHCO; = NaeCO3 + CO2 + HO 


A large part of the bicarbonate is converted to carbonate for 
the market by this process. The carbon dioxide obtained is, of 
course, returned to use in the first stage of the manufacture. 

Ammonia is much more valuable than sodium carbonate, at 
the present time, and the ammonia-soda process depends, 
economically, upon the recovery of the ammonia. ‘This is ef- 
fected by treating the solution of ammonium chloride with 
slaked lime, Ca(OH), and distilling: 


Z NH.,Cl + Ca(OH). — CaCl, +4 2 NH; 


In well-conducted factories not more than 5 kilos of ammonia 
are lost in preparing 1000 kilos of sodium carbonate. 

Sometimes magnesium oxide, MgO, is used in place of calcium 
hydroxide, as it is possible to recover chlorine or hydrochloric 
acid from the magnesium chloride, but the fact that the chlorine 
is left by the process in the form of a comparatively worthless 
compound may lead, ultimately, to its abandonment. 

Sodium bicarbonate is used as “ baking soda ”’ in cooking, to 
furnish carbon dioxide for lightening bread or cake. It is used 
with sour milk or, frequently, mixed with cream of tartar, tar- 
taric acid, alum or acid calcium phosphate in the various baking 
powders. It is also used as a mild alkali in medicine. 

Sodium Silicate or Soluble Glass, Na2SiOs, is made by fusing 
sodium carbonate and sand in the proper proportion. -It is 
hydrolyzed by water and the solution reacts strongly alkaline, 
but the silicic acid remains in colloidal solution. The solution 
is used to fireproof wood and fabrics, covering them with a thin, 
glassy coating, which renders them much less inflammable. 

Sodium Tetraborate or Borax, Na2B.07.10 H.O, has been con- 
sidered in connection with boron (p. 367). 


CHAPTER XXIV 


ALKALI METALS (Continued): POTASSIUM, AMMONIUM, 
RUBIDIUM, CASIUM: THE SPECTROSCOPE 


Potassium, K, 39.10. Occurrence. Many of the natural sili- 
cates and especially the potash feldspar, orthoclase, KAISi3QOs, 
contain potassium. These feldspars form an essential constitu- 
ent of granite and in the disintegration of granites and other 
rocks through the action of water during very long periods of 
geological time the sodium and potassium have been partly re- 
moved and carried away to the ocean. Owing to the selective 
action of the colloidal silicates remaining in the beds of clay and 
soils formed by the disintegration of the rocks, more potassium 
than sodium has been retained, and the element forms a constitu- 
ent of very great importance in all arable lands. From the soil, 
potassium is taken up by all plants in their growth and when 
vegetable material is burned, the ash is almost invariably 
alkaline from the presence in it of potassium carbonate. For- 
merly wood ashes were the most important source of potassium 
compounds. ‘The potassium carbonate was obtained from the 
ashes by leaching them with water, and the “lye” prepared in 
this way was used for the domestic preparation of soft soap. 
The latter is a concentrated solution of potassium salts of the 
organic acids of ordinary fats, such as lard or tallow. These 
salts are deliquescent and cannot be readily brought to a solid 
form, as is the case with the sodium salts, which form ordinary 
hard soap. 

While beds of common salt are found in many different parts 
of the world and strong brines are quite common, large deposits 
of potassium chloride (sylvite, KCl) and of magnesium potassium 


chloride (carnallite, MgCl,.KCI.6.H:;O) have thus far been 
414 


g 


POTASSIUM 415 


found only in Germany and especially at Stassfurt. These de- 
posits furnished, for a long time, the larger portion of the potas- 
sium compounds used in the world, and especially they furnished 
the potassium required to maintain the fertility of the soil for 
raising tobacco, cotton and other crops. 

During the Great War methods were developed for obtaining 
potassium compounds from the waters of Searle’s Lake in 
California, from the smoke’ of cement kilns and from other 
sources. Some of the new processes seem likely to withstand 
the competition of the imported products and we may hope to 
become independent of the foreign supply. 

Metallic Potassium was prepared first by Sir Humphry Davy 
in 1807 at about the same time that he discovered metallic so- 
dium. It may also be prepared by reducing potassium carbon- 
ate with carbon and by the electrolysis of fused potassium chlo- 
ride. Potassium is a silver-white metal which tarnishes instantly 
in moist air and takes fire when thrown on water. The hydrogen 
which is liberated burns with the characteristic violet flame of 
potassium. It has a specific gravity of 0.8621 at 20°. It melts 
at 62.3° and boils at about 760°. 

Potassium Oxide, K,O, has been prepared by the partial 
oxidation of metallic potassium in dry air followed by distilling 
away the excess of metal. It combines energetically with water 
to form potassium hydroxide, KOH. 

Potassium Hydroxide was formerly prepared by treating a 
solution of potassium carbonate with. slaked lime : 


K.COs + Ca(OH). = 4 KOH of CaCOsz 


The reaction depends, of course, on the relative insolubility 


of calcium carbonate. 

Potassium hydroxide is now prepared commercially from 
potassium chloride by electrolysis, with the Castner-Kellner 
and other forms of apparatus (p. 402). 

Potassium hydroxide is a white, deliquescent solid. The solu- 
tion in water has a soapy feel and attacks the skin strongly if 
allowed to remain in contact with it. The ordinary solid potas- 
sium hydroxide of the laboratory contains 15 to 20 per cent of 


416 A TEXTBOOK OF CHEMISTRY 


water. Sodium hydroxide, on the contrary, is usually almost 
anhydrous. 

Solutions of potassium hydroxide attack glass less than those 
of sodium hydroxide and do not give a precipitate of the carbon- 
ate as easily as the latter. For these reasons such solutions are 
usually employed in organic analysis and often in gas analysis, 
for the absorption of carbon dioxide. 

Potassium Chloride, KCI. The occurrence of potassium chlo- 
ride and of carnallite in Germany has been mentioned. It erys- 
tallizes in cubes and melts at about 750°. It is easily soluble in 
water. The solution is often used as a standard for electrical 
conductivity. 

The crude salt is extensively employed in fertilizers and a 
purer form is used in the manufacture of saltpeter, KNOs. 

Potassium Chlorate, KCIO3, may be prepared by saturating 
a warm solution of potassium hydroxide with chlorine (p. 127). 
Practically, milk of lime, Ca(OH)», is saturated with chlorine, 
forming calcium chloride and calcium chlorate, and to this solu- 
tion potassium chloride is added, causing the separation of 
potassium chlorate, which is not very easily soluble. The 
commercial reason for such a procedure is apparent. Potas- 
sium chlorate is also made by the electrolysis of a solution of 
potassium chloride under such conditions that the chlorine and 
potassium hydroxide formed react with each other. The final 
result may be expressed by the equation : 


Potassium chlorate is used in the preparation of oxygen, in 
medicine and in the manufacture of matches. 

Potassium Perchlorate, KClO,, is formed when potassium 
chlorate is heated slightly above its melting point : 


4 KCIO; = 3 KCIO, + KCl 


It is impossible to avoid some decomposition of the chlorate 
or perchlorate with evolution of oxygen, but with care a consid- 
erable portion of the chlorate may be converted into the per- 


POTASSIUM SALTS 417 


chlorate. The latter is much less soluble than the chloride or 
chlorate and may be purified by crystallization from hot 
water. 

Potassium Iodide, KI, may be prepared by dissolving iodine 
in a solution of potassium hydroxide. The iodate, KIO3, formed 
at the same time, may be decomposed by heating the mixture 
alone, or, better, with charcoal or some other reducing agent. 
The commercial salt often contains a little iodate, which is very 
objectionable for many laboratory uses. The salt is used in 
medicine and for the preparation of photographic plates. 

Potassium Polyiodides. Solutions of potassium iodide dis- 
solve iodine readily, forming unstable polyiodides, the one having 
thecomposition KI; being probably present in solutions of mod- 
erate concentration. These solutions dissociate easily into potas- 
sium iodide and iodine and react in the same manner as free 
iodine toward reducing agents. For this reason such solutions 
are used in volumetric analysis. The reaction of such a solution 
with sodium thiosulfate has been given (p. 187). 

Potassium Sulfate, K2SQOu,, crystallizes without water of crys- 
tallization. It melts at 1080°. It forms double salts with 
magnesium and calcium, which are found in the potash deposits 
in Germany and are an important source of potassium com- 
pounds. | 

Acid Potassium Sulfate, or Potassium Bisulfate, KHSOu., can 
be prepared by heating a mixture of potassium sulfate and sul- 
furie acid in molecular proportions. When heated gently it is 
converted into potassium pyrosulfate, K2S5207, with loss of water. 
At a higher temperature the latter loses sulfur trioxide and goes 
back to the normal sulfate. 

Potassium pyrosulfate is often used in the laboratory as a 
solvent for aluminium oxide, Al.O3, ferric oxide, Fe203, titanium 
oxide, TiOs, and other difficultly soluble substances. Sodium 
pyrosulfate is, however, more suitable for this purpose. (Hille- 
brand, Analysis of Silicate and Carbonate Rocks, p. 105.) 

Potassium Nitrate or Saltpeter, KNO;. After the introduc- 
tion of gunpowder into Europe, about 1300, and especially after 


418 A TEXTBOOK OF CHEMISTRY 


it came into general use in warfare in the sixteenth century, the 
preparation of pure saltpeter for use in its manufacture became 
continually more important. Until the nineteenth century 
saltpeter was obtained almost exclusively from natural sources, 
where it had been formed by the decay of organic matter con- 
taining nitrogen and potassium, in the presence of nitrifying 
bacteria. For a long time considerable supplies of saltpeter have 
been obtained from India, where it is formed in this way. 
Calcium nitrate, Ca(NO3)2, which is formed in a similar manner, 
sometimes occurs as an efflorescence on the walls of stables or in 
cellars. By interaction with potassium carbonate from wood 
ashes saltpeter is readily obtained. During the French Revolu- 
tion saltpeter was often prepared in this manner. During the 
War of 1812 the United States depended largely on saltpeter 
from the Mammoth Cave, Kentucky. 

After the discovery of Chili saltpeter, NaNOs, in South Amer- 
ica and of potassium chloride at Stassfurt, the manufacture of 
saltpeter from these salts was developed. The preparation de- 
pends on the fact that sodium chloride is about equally soluble 
in hot or cold water, while potassium nitrate dissolves in one 
half of its weight of water at 87°, but only 25 parts of the salt will 
dissolve in 100 parts of water at 15°. If potassium chloride is 
added in molecular proportions to a concentrated, hot solution of 
sodium nitrate, it will pass into solution, and sodium chloride, 
the least soluble of the four salts (NaNO3, KCI, NaCl, KNOs) 
present, will separate. From the mother liquors potassium ni- 
trate will separate on cooling, since that is the least soluble con- 
stituent in the cool solution. For the manufacture of gunpowder 
a salt entirely free from chloride must be prepared by recrys- 
tallization and centrifugal drainage. 

Potassium nitrate crystallizes inrhombic prisms. It melts at 
339°. It is used in the manufacture of gunpowder and in the 
curing of meats, especially of salt beef. It imparts to the meat 
a desirable reddish color. Taken in considerable quantities it 
is a poison. 

Gunpowder is a mixture of about 75 parts of saltpeter, 13 parts 


POTASSIUM SALTS 419 


of charcoal and 12 parts of sulfur. This corresponds very 
nearly to the equation : 


The explosion depends on the fact that the oxygen for burning 
the carbon is contained in the mixture and on the large volume of 
the nitrogen and carbon dioxide formed in the reaction. The 
heat of combustion also raises these gases to a high temperature, 
increasing the force of the explosion. In the burning of the 
gunpowder the grains burn from the surface inward, and the 
speed of the combustion is closely connected with the size of 
the grains. For use in large ordnance, hexagonal blocks an inch 
or more in diameter are used to secure slower combustion and 
allow time for the heavy shot to gain momentum before the full 
force of the explosive is developed. The failure of guncotton 
when it was first tried in firearms was partly due to the too rapid 
burning of the material, which caused the guns in which it was 
used to burst. This difficulty was finally overcome by giving 
the “‘ smokeless powder’ made from guncotton a dense form, 
somewhat resembling that of ordinary gunpowder. 

Potassium Nitrite, K.NQOz, is prepared by the reduction of the 
nitrate with lead, iron or sometimes with charcoal or sulfur. It 
is very easily soluble in water and is used in the laboratory as a 
reagent for cobalt, with which it forms a difficultly soluble com- 
plex salt, Co(NOz)3.3 KNOz or K3Co(NOz)s¢. The forma- 
tion of the same salt or of a similar salt containing silver, 
K,AgCo(NO:)s, may also be used as a test for detecting 
potassium. 

Potassium Carbonate, K2COs, was formerly obtained by leach- 
ing wood ashes. ‘The salt is also obtained in the scouring of wool 
and from the residual sirups of the beet sugar manufacture after 
the sugar of the sirups has been converted into alcohol. It is 
also made from potassium chloride by processes similar to those 
used in manufacturing sodium carbonate. 

Potassium carbonate is a deliquescent salt, differing in this 
respect very markedly from sodium carbonate. The anhydrous 


420 A TEXTBOOK OF CHEMISTRY 


salt melts at about 890°. It is used in making soft soap and hard 
glass. . 
Potassium Bicarbonate or Saleratus, KHCOs, is easily pre- 
pared by passing carbon dioxide into a concentrated solution 
of potassium carbonate. It was formerly used in cooking, but 
has been entirely displaced by the cheaper and more suitable 
sodium bicarbonate, NaHCO3. Potassium bicarbonate dis- 
solves in about three parts of water, being much more easily 
soluble than the sodium salt. The solution is nearly neutral 
to phenolphthalein, but loses carbon dioxide and becomes alka- 
line on boiling. 

Potassium Cyanide, KCN, can be prepared by heating potas- 
sium ferrocyanide : 


KyFeC.Ne¢ =a KCN 4- Fe oe 2 G + INy 


By heating the salt with metallic sodium a mixture of potas- 
sium and sodium cyanides is obtained and all of the: cyanogen 
of the original salt can be saved. This mixture may be used 
for nearly all purposes to which potassium cyanide is applied and 
especially for the extraction of gold from its ores. 

Ammonium, NH,, is a group which so closely resembles potas- 
sium in the salts which it forms with acid radicals that it seems 
desirable to speak of these salts at this point. It has not been 
found possible to separate ammonium, NH,, by itself, but if a 
solution of ammonium chloride is poured on some sodium 
amalgam the reaction represented by the equation : 


NH.Cl + Na(Hg) = NaCl + NH, (Hg) 


takes place and it can be shown by the electrical properties 
and by the effect in reducing metals from their salts that a small 
amount of ammonium amalgam is formed. ‘The substance is, 
however, extremely unstable and decomposes rapidly into am- 
monia, NH3, hydrogen and mercury. 

Ammonium Hydroxide, NH.OH. Solutions of ammonia in 
water seem to involve the following equilibria : 


NH3 + H20 iad NH.,OH = NH, + OH- 


AMMONIUM SALTS 421 


The ratio between these various substances has not been deter- 
mined with any degree of certainty. It is known, however, that 
the concentration of the hydroxide ions is small in comparison 
with the concentration in a solution of sodium or potassium 
hydroxide of equivalent molecular concentration. In other 
words, ammonium hydroxide is a comparatively weak base. 

Ammonium Chloride, NH,Cl, is prepared by neutralizing the 
aqueous gas liquors, obtained in the manufacture of coal gas or 
coke, with hydrochloric acid. It is partially purified by subli- 
mation. When heated to about 350° under ordinary conditions, 
it is converted into a gas which consists of a mixture of ammonia 
and hydrochloric acid, the weight of a gram molecular volume 
of the gas being only about 26.8 grams instead of 53.5 grams, 
as would be expected from the formula. This abnormal density 
was at one time used as an argument against Avogadro’s hy- 
pothesis. It has been shown by diffusion experiments, however, 
that the gas is a mixture, as the ammonia, which is the lighter 
of the two constituents, diffuses away more quickly than the 
hydrochloric acid. Finally, many years after the theory of dis- 
sociation had been universally accepted as the correct explana- 
tion of the abnormal density, it was shown that very carefully 
dried ammonium chloride may be converted into a vapor without 
dissociation, and that a gram molecular volume of this vapor 
weighs about 53.5 grams. 

Ammonium Sulfide, (NH,4)2S, is prepared by passing hydrogen 
sulfide, H.S, into a solution of ammonia. ‘The most convenient 
method is to take a known quantity of a 10 per cent solution of 
ammonia (sp. gr. 0.96) and pass into it the hydrogen sulfide 
generated by the action of one fifth of its volume of concentrated 
sulfuric acid upon an excess of ferrous sulfide, FeS, contained in a 
bottle holding ten volumes of water for one volume of the acid. 
The gas should be generated rapidly and well washed with water 
before entering the ammonia. 

Ammonium Hydrosulfide, NH.SH. If hydrogen sulfide in 
excess is passed into a solution of ammonia, the hydrosulfide is 
formed. When a solution of either ammonium sulfide or of 


4.32 A TEXTBOOK OF CHEMISTRY 


the hydrosulfide is exposed to the air, especially if exposed also 
to light, the sulfides are oxidized, sulfur separates and am- 
monium hydroxide is regenerated. The principal action is 
parallel to that of air on hydrogen sulfide water : 


Hes + O= H.O+S 


The sulfur liberated in this manner will, for a time, dis- 
solve in the ammonium sulfide, forming polysulfides, (NH4)2So, 
(NH4)0S3, ete. These polysulfides give to the solution a yellow 
color. After the oxidation has gone beyond a certain point, the 
separated sulfur no longer finds any ammonium sulfide with 
which to combine and begins to separate in the free state. The 
solution is then no longer fit for use as a reagent. 

Ammonium sulfide is used in the laboratory to precipitate 
those metals whose sulfides are too soluble for precipitation in 
acid solution but sufficiently insoluble for precipitation in the 
presence of a base. The polysulfide is used to convert stannous 
sulfide, SnS, into stannic sulfide, SnSs, and to dissolve the latter 
in separating it from lead, bismuth and other metals. Either 
may be used to dissolve arsenious sulfide, As2Ss, or antimony 
sulfide, SbS3 (p. 261): 

Ammonium Sulfate, (NH4).SOu, is prepared in a crude form 
by neutralizing the ammoniacal liquors of the gas works with 
sulfuric acid. As sulfuric acid is the cheapest of the commercial 
acids, this salt is often prepared to put the ammonia into suit-— 
able form for transportation or for use in fertilizers. Ammonia 
can, of course, be readily regenerated from it by treatment with 
lime, CaO. 

Ammonium Nitrate, NHiNOs, may be prepared by neutraliz- 
ing nitric acid with ammonia or ammonium carbonate and 
evaporating the solution to crystallization. It is easily soluble 
in water. At 166° it melts and decomposes into nitrous oxide, 
N2O, and water. The decomposition is exothermic, and if the 


AMMONIUM SALTS 423 


temperature is too high or if a large amount of the salt is heated 
at once, the reaction may become explosive : 


NH,NOs3 = NO + ye H.O -++ 29,900 cal. 


The salt is also used as a very important constituent of modern 
explosives. | 

Ammonium Nitrite, NHsNOz, is a deliquescent, very unstable 
salt, which decomposes easily into nitrogen and water. 

Ammonium Sodium Hydrogen Phosphate, or microcosmic salt, 
NaNH,HPO,.4H20, is used in blowpipe analysis to furnish a 
bead of sodium metaphosphate, NaPO;, which, when hot, will 
dissolve many metallic oxides, giving characteristic colors. The 
bead does not dissolve silica, SiOz. 

Ammonium Carbonate, (NH,)2CO3. A salt known commer- 


cially as “‘ammonium carbonate ” 


is prepared by heating a 
mixture of calcium carbonate and ammonium sulfate. It con- 


sists of a mixture of ammonium bicarbonate, ,.NHsHCOs, and 


O 
ammonium carbamate, NH,—O—CZNH,. The second of 
these salts may be prepared, also, by the direct union of carbon 
dioxide and ammonia. By dissolving the commercial carbonate 
in water and adding ammonia, the bicarbonate, NHsHCQs, is 
changed to the normal carbonate, (NH:),CO3. The carbamate 
is also soon hydrolyzed to the normal carbonate : 
O O 
NH,—O—CZNH, +H,0= NH,—O—CZ0—NH, 

A solution prepared in this way is used as the ordinary labora- 
tory reagent. 

Ammonium Bicarbonate, NH:HCOs, is prepared by passing 
carbon dioxide into a solution of ammonia as one of the opera- 
tions of the ammonia-soda process (p. 412). 

Ammonium Chloroplatinate, (NH,)2PtCls, is a very difficultly 
soluble salt closely resembling the corresponding potassium salt. 
Similar compounds are formed from many amines, compounds 
in which one or more hydrogen atoms of ammonia have been re- 
placed by organic radicals. 


424 A TEXTBOOK OF CHEMISTRY 


Rubidium, Rb, 85.45, and Cesium, Cs, 132.81. In 1860 and 
1861 Bunsen and Kirchoff in applying their newly discovered 
method of spectrum analysis to the study of a mineral water 
from Durkheim found some spectral lines which did not corre- . 
spond to those of any known element. In order to obtain enough 
material to study the compounds of the new elements, 40 tons 
of the water were evaporated and the compounds of rubidium 
and caesium were extracted from the residues. Rubidium was 
named from the Latin word rubidus, meaning red, and cesium 
from the Latin cesius, the blue of the sky, because of the red 
lines in the spectrum of the former and the blue lines in the 
spectrum of the latter. Rubidium is found to the amount of 
about 0.025 per cent in carnallite (KMgCl3.6H2O) and as a 
million and a half tons of this mineral are worked over annually 
for the preparation of potassium compounds it would be possible 
to obtain very considerable quantities of the element. 

Metallic rubidium melts at 38.5°, csesium at 26.5°, the lowest 
melting point of any metal except mercury. Their compounds 
resemble those of potassium in their general properties. The 
chloroplatinates, RbePtCle, and Cs2PtCl, and the alums, 
RbAI(SO4)2.12 HO and CsAl(SO,)2.12 H:O, are less soluble 
than the corresponding potassium compounds and are used in sep- 
arating the elements. A chloroiodide of ceesium, CsClel, is also 
especially useful for the preparation of pure cesium compounds. 

Spectrum Analysis. arly in the nineteenth century Fraun- 
hofer pointed out that when a solar spectrum is produced in such 
a manner that the colors are sharply separated from each other, 
the spectrum is crossed by a series of dark lines. During the 
years from 1820 to 1860 several different observers noticed the 
characteristic colors imparted to flames by different elements and 
the spectra of bright lines given by these colored flames and also 
those given by metals which are vaporized by the electric spark. 
It was not till 1860, however, that all of these phenomena were 
brought into clear relationship by the classical researches of 
Kirchoff and Bunsen, which culminated in the discovery of 
rubidium and ceesium. 


THE SPECTROSCOPE 425 


The simplest form of spectroscope is shown in Fig. 97. The 
prism A is of glass having a high dispersive power. A narrow 
slit at B is illuminated by the flame or light which is to be exam- 
ined. Between the slit and the prism is placed a lens at C, 
which renders the rays of light from the slit parallel before they 





Fig. 97 


reach the prism. ‘The prism is set in such a manner that the 
angle of incidence on the first face is the same as the angle of 
emergence from thesecond face, as this gives the purest spectrum. 
The light is examined by means of the telescope D. A scale 
placed at EH, illuminated by a light placed before it and whose 
image is reflected from the surface of the prism, serves to locate 
the position of the lines. 

Such a flame as that of acetylene gives a continuous spec- 
trum, indicating that molecules of solid carbon in the white 
flame are vibrating in all possible periods required to give white 
light. Ifa Bunsen flame is placed before the slit and some com- 
pound of sodium, as sodium chloride, is introduced, the flame 
assumes a brilliant yellow color, and with a single prism spec- 
troscope the spectrum consists of a single, bright yellow line. 


426 A TEXTBOOK OF CHEMISTRY 


A spectroscope having several prisms, or a spectroscope using 
a metallic mirror (‘ grating ’’) ruled with many thousands of 
equidistant lines and which gives a diffraction spectrum, will 
separate the line into two lines situated close together. The 
wave lengths of the lines are 0.5896 and 0.5890 microns, the 
micron being the thousandth part of a millimeter. The physi- 
cal significance of these lines seems to be that under the con- 
ditions of the flame either the sodium atoms as a whole or, 
more probably, portions of the sodium atoms or electrons within 
or around them vibrate at a definite rate, which is independent 
of the temperature. This rate is almost inconceivably rapid. 
The velocity of light is about 300,000 kilometers per second. 
This is equal to 3 X 10 microns, and since the wave length of 
the sodium light is only 0.59 micron, the number of vibrations 


Ee SA ee ==—5) 4 10"4 per 


per second must be approximately REG 


second. 

There are two dark lines in the solar spectrum which coincide 
exactly with the bright yellow lines of the sodium spectrum. 
This is explained by supposing the interior of the sun to be an 
incandescent mass which gives out light vibrations of all wave 
lengths corresponding to the visible spectrum. The photo- 
sphere of the sun, on the other hand, consists of a gaseous en- 
velope or atmosphere containing many different elements, among 
these sodium. The sodium atoms, if they have the power of 
producing light waves in the ether by their vibrations, must also 
be able to absorb waves of the same length from the ether, 
exactly as a tuning fork is set in vibration by sound waves of its 
own pitch, while waves of a different pitch do not affect it. The 
sodium atoms in the photosphere, therefore, absorb the waves 
of their own particular rate; and while they give the energy ab- 
sorbed back again to the ether, they dissipate the energy by 
spreading it in all directions instead of allowing it to pass on 
toward the observer. The result is that the portion of the 
spectrum corresponding to the sodium vibrations will be rela- 
tively dark. By means of this principle it has been possible to 


THE SPECTROSCOPE Ay, 


show that more than thirty elements found on the earth are found 
also in the sun. One of these elements (helium), indeed, was 
discovered in the sun before it was found on the earth. 

For the purpose of comparing spectra it is convenient to place 
a right-angled prism before the slit of the spectroscope in such a 
manner as to cover one half of it. This may be made to reflect 
the light from a second flame into the slit in such a way that 
the spectrum from one flame will occupy the upper half of the 
field of vision while the spectrum from the other flame will occupy 
the lower half. In this manner the coincidence of lines in the 
two spectra may be readily observed. 

In another form, known as the direct vision spectroscope, a ° 
series of prisms of different kinds of glass are so combined that 
one kind of glass counterbalances the mean refractive index of the 
_ other, while the dispersive effects are not counterbalanced. 
The effect is exactly the reverse of that in an achromatic lens. 
Such spectroscopes are especially suitable for the detection of the 
alkali and alkali-earth metals in qualitative analysis. By means 
of the spectroscope it is possible to detect zya007 milligram 
of sodium. Only the methods used in studying radioactive sub- 
stances are more sensitive than this. 

To obtain the spectra of iron, copper and other metals, which 
are volatile only at high temperatures, electric sparks from a 
Rumkhorf coil are passed between terminals of the metal, or, 
in some cases, between platinum wires, one of which is in a small 
cup containing a solution of a salt of the metal. The spectra 
of gases are observed in Plucker tubes, which have a narrow 
portion through which the electric discharge is passed. 


CHAPTER XXV 


THE ALTERNATE METALS OF GROUP I. COPPER, 
SILVER, GOLD. PHOTOGRAPHY 


CopPER, silver and gold, which alternate with potassium, rubid- 
ium and cesium in Group I of the Periodic System, are in 
almost the greatest possible contrast with those metals. The 
alkali metals are light. They melt and volatilize at compara- 
tively low temperatures and they react violently with water at 
ordinary temperatures. Copper, silver and gold are heavy 
metals, all three melt between 960° and 1083° and none of them 
decomposes water, even at high temperatures. They are also 
the best conductors of electricity that we have. As they do 
not decompose water, all three of these metals are found free 
in nature. All three have been known and used since very 
early times. 

Copper, Cu, 63.57. Occurrence. Copper is found free in 
nature, especially in the Lake Superior region. 

It is found as copper pyrites, or chalcopyrite, CuFeSe, a mineral 
closely resembling iron pyrites in superficial appearance but 
having a different crystalline form and usually showing blue, 
red or green colors, owing to superficial changes. Bornite, 
CusFeS3, chalcocite, cuprous sulfide, CuzS, and malachite, 
Oa 
CuCO3.Cu(OH)s, or CO 


€ , a basic carbonate, are 
O—Cu—OH 
the other most important minerals containing copper. 
Metallurgy. The most common ores of copper contain either 
the sulfide, CusS, or copper pyrites, CuFeS,. If the ore is 
poor in copper, it is sometimes concentrated by crushing it 
and washing away a part of the lighter minerals with water. 


The concentrated ore is then roasted in a furnace with the 
428 


COPPER 429 


addition of sand, if enough silica is not already present. The 
iron is partly oxidized to ferrous oxide, FeO, which combines 
with the silica, SiOz, to form a fusible silicate or slag, FesSiOu. 
The cuprous sulfide and some of the ferrous sulfide melt and 
sink to the bottom of the furnace beneath the slag, which is 
much lighter. The mixture of sulfides obtained in this way is 
drawn off and is known as copper matte. Similar operations are 
now often carried out in a blast furnace somewhat similar to 
that used in the manufacture of pig iron (p. 543). In the older 
processes the ore was usually carried through a long series of 
complicated operations for the purpose of securing a compara- 
tively pure matte and reducing the latter to metallic copper. 
In the United States these processes have been very much 
shortened by the use of a modification of the Bessemer Converter 
(p. 547) for the reduction of the matte. The molten matte 
from the roasting furnace is poured directly into the converter, 
where it is subjected to a blast of air mixed with fine sand or 
silica. The sulfur of the ferrous sulfide and cuprous sulfide is 
burned out, the heat of the combustion maintaining the tem- 
perature of the converter. The ferrous oxide combines with 
the sand to form a slag of ferrous silicate, while the copper 
melts and may be cast into plates. The principal reactions 
may be expressed as follows: 


2FeS + 30, = 2 FeO + 2S0, 
2 FeO a SiO. = Fe SiO, 
Cues aa 202 == CuO + SO. 
2d Stage 2 CuO + Cues = 4Cu + SO, 
or Cues + O2 = 2Cu + SO2 


Ist Stage | 


Electrolytic Refining of Copper. The copper obtained by 
means of the Bessemer converter or by any of the more com- 
plicated furnace methods usually contains a small amount of 
gold and silver and larger quantities of arsenic, lead and other 
metals which render it unfit for most industrial uses and es- 
pecially for use in electrical conductors. Nearly all of the 
copper. is now refined electrolytically. The plates of crude 


430 A TEXTBOOK OF CHEMISTRY 


copper are suspended, upright, in a long tank (Fig. 98) filled 
with a solution of copper sulfate. By connecting the two end 
plates with the poles of a dynamo the current flowing through 
the system will cause each plate to become negative on one side 
and positive on the other. On the positive side the copper will 


a = 





Fig. 98 


pass into solution as cupric ion, Cut*. Gold, silver, bismuth 
and some other impurities in the copper fail to dissolve and are 
collected as ‘‘slimes”’ and treated for the recovery of silver and 
gold. On the negative side of each plate nearly pure copper 
will be deposited. When all of the original copper has been 
dissolved the plates of pure copper are removed and new plates 
of crude copper put in their place. 

The production of copper in the United States in 1910 was 
900,000 tons, worth $137,000,000. 

Properties of Copper. Copper is red as ordinarily seen by 
reflected light. In very thin films it transmits green light. It 
melts at 1083° (1063° in air, because of oxidation), boils at 2310° 
and has a density of 8.93. It is the best conductor of electricity 
of the cheaper metals. Its conductance is very much impaired, 
however, by the presence of small amounts of other metals. 
Arsenic is especially harmful, 0.03 per cent lowering the con- 
ductance by 14 per cent. <A table of conductances for the more 
common metals will be found at the close of this chapter. 

When exposed to the weather, copper slowly covers itself 


ALLOYS OF COPPER 431 


with a green coating of basic carbonate of the composition of 
malachite, CuCO3.Cu(OH)2. When heated in the air a coating 
of the black oxide, CuO, is formed, which comes off in scales 
on cooling. In the absence of air, copper is not affected by 
hydrochloric or dilute sulfuric acid. It dissolves in nitric acid 
as copper nitrate, Cu(NOs)s, with evolution of nitric oxide, NO, 
nitrous oxide, NO, nitrogen peroxide, NOs, or a mixture of 
these according to the concentration of the acid and the tem- 
perature. Hot, concentrated sulfuric acid dissolves it as cupric 
sulfate, CuSO,, while sulfur dioxide is evolved. 

Copper is used for electrical conductors, for the sheathing of 
ships and for much of the apparatus used in the fermentation 
industries. 

Alloys of Copper. Copper is used in a great variety of alloys, 
the two most important being brass and bronze. Brass con- 
tains 60 to 70 per cent of copper and 40 to 30 per cent of zine, 
with usually small amounts of lead, tin and iron. Bronze is 
primarily an alloy of copper and tin, usually with 80 to 90 per 
cent of copper; it contains, in most cases, small amounts of zinc 
and lead. Bronze is used for bells and statuary and was for- 
merly also used for cannon. Phosphor bronze contains from 
0.25 to 2.5 per cent of phosphorus, which makes it hard and 
suitable for bearings in machinery. It has recently been dis- 
covered that the addition of a very small amount of copper to 
iron or steel greatly increases its resistance to corrosion. 

Copper Hydroxide, Cu(OH)p, is precipitated as a blue, amor- 
phous compound when a solution of sodium hydroxide is added 
to a solution of copper sulfate or any similar salt. If the solu- 
tion is heated to boiling, the hydroxide is decomposed into the 
black cupric oxide and water : 


Cu(OH), = CuO + H,O 


The instability of copper hydroxide is in very marked con- 
trast with the stability of the hydroxides of the alkali metals, 
but the hydroxides of silver and mercury are still less stable 
and have never been prepared as pure compounds. 


432 A TEXTBOOK OF CHEMISTRY 


Cupric Oxide, CuO, may be prepared, as just described, by 
the decomposition of the hydroxide, or by the decomposition 
of the nitrate, but neither method will give a pure compound. 
The copper oxide which is used for organic analysis is usually 
prepared by calcining copper wire in the air for a long time. 
Copper oxide obtained by heating the nitrate retains nitrogen 
obstinately and is not suitable for use in the determination of 
nitrogen in organic compounds. 

Cuprous Oxide, Cu2O, is formed as a bright red precipitate 
by the action of reducing agents and especially by the action 
of hydroxylamine, hydrazine, glucose, fructose or lactose on a 
solution of copper sulfate, sodium hydroxide and potassium 
sodium tartrate (p. 335). 

Cupric Chloride, CuCl».2H;0, is a bright green salt obtained 
by dissolving cupric oxide in hydrochloric acid. The double 
salt, CuCl,e.2KCI.2 H,.O, is used to dissolve iron or steel without 
evolution of gas for the determination of carbon. 

Copper chloride dissolves in a small amount of water to a 
green solution similar in color to the crystallized salt. On 
dilution the color changes to the blue color characteristic of all 
solutions containing the cupric ion, Cutt. If not too dilute, 
concentrated hydrochloric acid will restore the green color, 
because the excess of chloride ions causes a decrease in the 
ionization of the cupric chloride. 

Cuprous Chloride, Cu2Cle, may be easily prepared by digesting 
a solution of cupric chloride, CuCle, in concentrated hydro- 
chloric acid with copper turnings : 


CuCl, a Cu = CueCl, 


Cuprous chloride is very difficultly soluble in water but dis- 
solves rather easily in concentrated hydrochloric acid, owing to 
the formation of chlorocuprous acid, HCuegCls. On dilution 
this is dissociated into its components and cuprous chloride 
is precipitated. The insolubility of cuprous chloride, CueCls, is 
parallel to that of silver chloride, AgCl [Ag2Cl.], and mercurous 
chloride, HgeCle. 


COPPER SALTS 433 


Cuprous chloride dissolves easily in ammonia, also. It is 
supposed that a complex molecule of the form, [Cu(NHs;)n]Cl, 
is formed, but the exact composition of this molecule has not 
been determined. 

Solutions of cuprous chloride either in hydrochloric acid or 
in ammonia absorb carbon monoxide and are used for that 
purpose in gas analysis. 

Sodium hydroxide decomposes cuprous chloride with the for- 
mation of cuprous oxide. Cuprous hydroxide, CuOH, has not 
been prepared as a definite compound, and there is some reason 
to suppose that it could exist only at low temperatures. 

* Cuprous Iodide, CuzlIz. If potassium iodide is added to an 
acid solution of a cupric salt, cuprous iodide is precipitated and 
iodine is liberated : 

2 Cu(NOs3)2 as 4 KI = Cul, + 4 KNO; ao I, 

The reaction is quantitative, and by titrating the iodine with 
a solution of sodium thiosulfate the amount of copper present 
may be determined. Cuprous iodide is white and nearly in- 
soluble in water. 

Cupric Sulfide, CuS, is formed as a black precipitate on pass- 
ing hydrogen sulfide into an acid solution of a cupric salt. It 
is very insoluble in water or dilute acids, but dissolves readily 
in warm nitric acid, the sulfur separating mostly in the free 
state. On ignition in a current of hydrogen it is converted into 
cuprous sulfide, Cuz2S, which has the same composition as the 
mineral chalcocite. 

Copper Sulfate, or Blue Vitriol, CuSO4.5H.O, forms deep blue 
crystals of the triclinic system. At 120°-140° these lose four 
fifths of their water, leaving the hydrate, Tae H.O. This 


YONI 
has, very probably, the structure Cuc >séon, At 240° 
O 


the last molecule of water may be expelled, leaving the anhydrous 
salt, CuSO, as a white powder. Copper sulfate is the most 
common salt of copper. It is used in the electrolytic refining 
of copper, in electroplating and electrotyping, as a mordant in 


434 A TEXTBOOK OF CHEMISTRY 


dyeing and in the gravity cells formerly much used for tele- 
graphic purposes. The anhydrous salt is sometimes used for 
drying alcohol, as it is very hygroscopic. 

Vitriols. Vitriol is a very old name given to sulfates because 
many of them have a glassy appearance. It is rarely used, now, 
except for copper sulfate, CuSO.4.5H2O, or blue vitriol, ferrous 
sulfate, FeSO,.7 H.O, or green vitriol, zine sulfate, ZnSO,4.7H.O, 
or white vitriol, and for sulfuric acid, oil of vitriol. The last 
name is derived from an old method of preparing the acid by 
distilling green vitriol. | 

* Cupric Nitrate, Cu(NO3):..6H.O, is formed in blue tabular 
crystals when solutions of the salt are crystallized at tempera- 
tures below 24.5°. Above that temperature the hydrate 
Cu(NOs3)2.3H2O is obtained. On heating moderately the salt 
loses water, and at a higher temperature it is decomposed into 
copper oxide, CuO, oxygen and nitrogen dioxide, NOx. 

Ammoniocupric Sulfate, CuSO,.H2O.4NH3;. When ammonia 
is added to a solution of copper sulfate, a very intense blue 
color is produced. Alcohol precipitates from such a solution 
ammoniocupric sulfate. This may be considered as the hydrate 
of copper sulfate in which four molecules of water have been 
replaced by four molecules of ammonia. Similar compounds 
are formed with other salts of copper, and the reaction may be 
used for the detection of minute quantities of the element. The 
formula, [Cu(NHs3)4|SO4.H2O, used by Werner and others brings 
out more clearly the intimate relation between the ammonia 
and the copper. 

*Cuprous Cyanide, Cus(CN)2. Solutions of copper salts react 
with a hot solution of potassium cyanide in very much the 
same manner as with potassium iodide: 


2 CuSO, + 4 KCN = Cu2(CN)e + 2 K2S04 + C2Ne 


If the compounds are used in the proportions given in the 
equation, the cuprous cyanide separates as a white precipitate. 
If an excess of the potassium cyanide is used, a complex salt, 
K;Cu(CN),, which is easily soluble, is formed. This salt gives 


KLECTROMOTIVE SERIES 435 


a solution containing so few copper ions that no precipitate is 
formed by hydrogen sulfide. This property is sometimes used 
to separate copper from ‘cadmium, as the latter is precipitated 
as cadmium sulfide, CdS, under similar conditions. 

Precipitation of Copper by Iron, Electromotive Series. If an 
iron nail is dipped in a solution of copper sulfate or copper 
chloride, metallic copper will be deposited and iron will pass 
into solution. We may formulate the reaction thus: 


Cutt + Power + + Cu 


The copper ion, Cu**, loses its electrical charge, giving it to the 
iron, while the copper assumes the metallic form. If a strip of 
copper and one of iron are suspended in any ordinary electrolyte, 
an electrical current will pass from the copper to the iron on 
connecting the two metals by means of a wire, while within the 
electrolyte the positive ions will travel toward the copper and 
the negative ions toward the iron. As before, the iron will 
pass into solution and copper will be deposited on the copper 
strip, if the electrolyte is copper sulfate. On the basis of similar 
experiments a table of the metals may, be arranged in such a 
manner that each metal in the series will be positive toward all 
of the metals on one side of it and negative toward those on the 
other side. 

The simplest method of looking at these phenomena is to 
consider that each metal in contact with a solution has a cer- 
tain solution-pressure which tends to cause the metal to pass 
into solution. We may suppose that a few atoms do actually 
leave the piece of metal and pass into solution as ions, but this 
would give the solution a positive charge and leave the metal 
negative; and unless some means is provided for the escape of 
the electrical charges the electromotive force set up will very 
quickly balance the solution pressure and the process will cease. 
If the metal is dipped in a solution of one of its salts, the ions 
of the metal in solution will partly or wholly counterbalance 
the solution pressure of the metal, the quantitative effect 
depending on the concentration of the ions and the nature of the 


436 A TEXTBOOK OF CHEMISTRY 


metal. For many metals this effect may exceed the solution 
pressure of the metal and the latter will assume a positive poten- 
tial with reference to the solution through the deposit and dis- 
charge of some metal ions on the plate. 


SINGLE ELECTRODE POTENTIALS OF THE ELEMENTS AT 25° 
IN SoLuTIONS Havine AN “Activity” or UNITY 


ELECTRODE ELECTRODE REACTION Be 

Piel Lis 46+ E> 2.958 
Rinehp Rb = Rbt 4+ E- 2.924 
Keke K = Kt + E- 2.922 
Na, Nat Na = Nat + E- i he 
Ba, Batt Ba = Bat*+ + 2 E- 1.838 
he Ae) Sra ort +2 Ke 1.791 
Hp, OH- (Pt) 5H» 4. OH- = H.O > E- 0.828 
Zn, Zntt Zn = Zn** +2 E- 0.758 
Fe, Fet* Fe = Fet+ +2 E- 0.44 

Cd, Cdt* Cd = Cdt* +2 E- 0.398 
A Ped Eg Tl = Tl? + E- 0.336 
Sn, on? Sn = Sntt +2 E- 0.136 
Pb. Pb Pb = Pb*t + 2 E- 0.122 
Ho Hee) 4H, = H+ + E- 0.000 
Cu, Cutt Cu. = Cutt 4+ 2 E- — 0.345 
Hg, Hge*t 2He = Hes +2 KT — 0.799 
Ag, Agt g = Agt+ E- — 0.800 
Ins (Eb I- = 43L+E- — 0.536 
Bre BroPt) Br- = 4Br, + E- — 1.066 
Chel atry = Cl, + E- — 1.359 


* The above table! gives the (hypothetical) difference in 
potential between the metal, or between platinum covered with 
hydrogen or a halogen, and a solution containing one mol of the 
ion of the element in one liter. The difference in potential 
between platinum covered with hydrogen and a solution contain- 
ing one mol of hydrogen ions in one liter is assumed, arbitrarily, 
to be zero. The actual measurements on which the table is 
based were made with solutions of a different concentration and 
were calculated back to molal concentrations very much as we 
calculate the weight of a gas for standard conditions. Although 
the electrode potentials of copper, mercury and silver are given 
negative values in\the tabie, they are called electropositive ele- 


1 R. H. Gerke, Chemical Reviews, 7, 381 (1925). See also the 
International Critical Tables, edited by E. W. Washburn and pub- 
lished by the McGraw-Hill Book Co. 


ELECTROMOTIVE SERIES 437 
ments because their atoms lose electrons when they pass into 
solution. Chlorine, bromine and iodine, whose atoms take up 
electrons in passing into solution, are electronegative elements. 

This table may be used in two ways: first, any metal of the 
table may be precipitated by any other metal which has a 
lower absolute potential and it will = 
precipitate any metal with a higher 
potential — thus copper will pre- 
cipitate silver, but it will be precip- 
itated by lead or iron; second, the 
values may be used to calculate 
the electromotive force of a galvanic 
cell in which two of the metals are 
used. For an accurate calculation 
it 1s necessary to take account of the 
concentration of the solution in con- 
tact with each electrode, and when 










Pe ry RY 
BARES 
A c 
J 


two solutions are used, the difference eve on 
e 6 e Dy eS A 
in potential between the solutions SEAT PES 





must also be considered. Fic. 99 

The common gravity cell (Fig. 99) 
consists of a copper plate A, in contact with a solution of copper 
sulfate, and a zinc plate B, in contact with a solution of zinc sul- 
fate. The electromotive force of the battery is approximately 
the difference between the electrode potentials of copper and 
zinc, or + 0.34 — ( —0.76) = 1.10 volts. 

The electromotive force of the Weston cells, consisting of 
mercury in contact with mercurous sulfate and cadmium in 
contact with cadmium sulfate, would be for normal solutions of 
each + 0.80 — (— 0.40) = + 1.20. 

The Clark cell, which has zine in place of the cadmium, 
would have for normal solutions an electromotive force of 
+ 0.80 — (— 0.76) = + 1.56. In both cases, however, the 
electropositive metal (cadmium or zinc) is in contact with a 
concentrated or saturated solution of its salt, and the mercurous 
sulfate is only very slightly soluble. This lessens the difference 
of potential, and the actual value for the Clark cell is 1.422 


438 A TEXTBOOK OF CHEMISTRY 


volts, and for the Weston cell, in which a saturated solution of 
cadmium sulfate is used, it is 1.018 volts at 20°. 

Faraday’s Law. If the same electrical current is passed 
through a series of cells containing electrolytes which are de- 
composed by the current, it is found that the amounts of the 
elements which separate at the electrodes will be proportional 
to the equivalents of the elements, 7.e. to the atomic weights of 
the elements divided by their valences. In a series of cells con- 
taining solutions of the following compounds : 


H.SO, AgNOs Gul CuCl, SnCl, SnCl, 
HH, O07 Ase 0 Ga .Gi Cipencl pnG) sn Cl 
lg8g 108e¢8¢ 63.6¢35.5¢ 318¢35.5¢ 59¢35.5¢ 29.5¢35.5¢ 


the same current which causes the liberation of one gram of 
hydrogen will liberate 8 grams of oxygen, 35.5 grams of chlorine, 
108 grams of silver, 63.6 grams of copper from cuprous chloride, 
31.8 grams of copper from cupric chloride, 59 grams of tin from 
stannous chloride and 29.5 grams of tin from stannic chloride. 
This relation was discovered by Faraday in 1834 and is known 
as “ Faraday’s Law.” The rational explanation of the law is 
that each univalent ion is associated with a unit charge of elec- 
tricity, each bivalent ion with twice this unit charge and a 
quadrivalent ion with four times this amount. In accordance 
with the electron theory this unit charge is the charge carried 
by an electron, and the fact that it does not seem possible to 
account for Faraday’s law without assuming that there is a 
definite, unit charge, is one of the most important facts sup- 
porting the theory. 

It should not be overlooked that the amount of energy re- 
quired to decompose a gram equivalent of different electrolytes 
is not the same. The difference in potential between the two 
electrodes varies from cell to cell; and the energy consumed in 
the cell depends on two factors — the quantity of electricity 
passing through the cell and the fall in potential from one elec- 
trode to the other. The latter depends on the absolute poten- 
tials of the substances separating at the two electrodes and on 


SILVER 439 


the concentration of the electrolyte exactly as in the galvanic 
cells discussed in the preceding paragraph. An electrolytic cell 
may be considered as a galvanic battery in which the direction 
of the current has been reversed by the application of an external 
electromotive force. 

Silver, Ag, 107.88. Copper, silver, gold and all metals which 
are more electronegative than hydrogen in the electromotive 
series are found sometimes in the free state in nature. The 
only other metals found in the free state are iron, cobalt and 
nickel, and possibly some others which have come to the earth 
in the form of meteorites and have not had time since their 
arrival to become completely oxidized. 

Silver is also found as the sulfide, AgeS, either alone, or more 
often associated with other sulfides and especially with lead 
sulfide or galena, PbS. It is also found as the chloride in the 
mineral cerargyrite, or horn silver, AgCl. 

Metallurgy. In the electrolytic refining of copper a chloride 
is added to the electrolyte, and this causes the silver to separate 
as the insoluble chloride, AgCl, with the slimes from which the 
silver and gold are recovered. 

The metallic lead obtained from galena, PbS, always contains 
some silver. This is recovered either by Pattison’s process of 
crystallization or by Parke’s process of extraction with zinc. 

* Pattinson’s Process depends on the principle that a solution 
melts at a lower temperature than the solvent. If lead contain- 
ing silver is melted in an iron pot, on cooling, crystals of nearly 
pure lead separate at first, leaving a solution or alloy of silver 
and lead which is richer in silver than the original metal. The 
crystals of lead are skimmed out and transferred to another pot 
on one side while the richer alloy is transferred to a pot on the 
opposite side. By repeating the operation several times, nearly 
pure lead is obtained at one end of the series of pots and an 
alloy comparatively rich in silver at the other end. This rich 
alloy is then heated in a furnace with free access of air till the 
lead is oxidized to litharge, PbO, leaving very nearly pure 
silver behind. While this process is no longer used, it is of 


440 A TEXTBOOK OF CHEMISTRY 


historical interest and also interesting because of the principles 
of crystallization involved. 

Cupellation, Assaying. The process of oxidizing lead contain- 
ing silver is often carried out on a small scale in a muffle furnace 
(Fig. 100) on a small cup of porous bone ash, called a cupel, 
which absorbs the 
litharge. Small 
amounts of other 
metals are oxi- 
dized with the 
lead and absorbed 
by the cupel so 
that an almost 
pure silver bead 
remains. This 
can be weighed 
as a means of 
determining the 
amount of silver 
in lead _ bullion. 
The gold and sil- 
ver of an ore may 
be concentrated 
in a lead button 
which can then be 
cupelled. Such 





determinations, 
or ‘“‘ assays,”’ fur-— 
/ nish the basis for 
~ commercial trans- 
i) ti ith 
Wy, YY} ac eae ores 
MUFFLE FURNACE and bullion. 
Fig. 100 Parke’s Proc- 


ess for the ex- 
traction of silver from lead is exactly analogous to the extraction 
of a substance from water by means of ether. When zinc and 


SILVER . 44] 


lead are melted together, on allowing the mixture to stand for 
a few minutes in the melted condition, nearly all of the zinc 
will rise to the top in the form of a solution of lead in zinc, 
containing only a small quantity of lead. The solution of 
zine in lead below will contain only a small amount of zinc. 
The silver and gold, however, are much more soluble in zine 
than in lead, and a comparatively small amount of zinc will 
carry nearly all of the silver to the top with it. If much silver 
is present, it may be desirable to repeat the process. The 
alloy of zine and silver, which is skimmed from the top, may 
be placed in a retort and the zine distilled away. The residual 
lead is then removed by cupellation. 

* Amalgamation Process. From ores containing little lead or 
copper, silver was formerly often separated by an amalgamation 
process. The ore was pulverized and intimately mixed with 
metallic mercury and water. The mercury dissolved, or amal- 
gamated with the silver, if that was in the free state. If present 
as the chloride, the mercury reduced it to the metallic state : 

AgCl + Hg = Ag + HegCl 
Mercurous 
Chloride 

If the silver was present as the sulfide, it might be reduced by 
adding iron turnings, or the ore was roasted with salt before it 
was amalgamated : 


2 NaCl + AgeS = NaS + 2 AgCl 


After the amalgamation was complete the lighter materials of 
the ore were washed away from the amalgam in a current of 
water, the amalgam was collected, and the larger portion of the 
mercury removed by squeezing it through chamois skin, which 
allows mercury but not the solid amalgam to pass. The latter 
was heated in a retort to remove the rest of the mercury. 

Other Processes for the Recovery of Silver. The recovery of 
silver in the electrolytic refining of copper has been referred to 
above. The cyanide process (p. 446) has largely displaced all 
others for obtaining both silver and gold. 


442 A TEXTBOOK OF CHEMISTRY 


The production of silver in the United States in 1910 was 
57,000,000 troy ounces, worth $31,000,000. 

Properties of Silver. Alloys. Silver is a soft, white metal, 
extremely malleable and ductile. It is the best conductor of 
electricity known. (See table at the end of this chapter.) It 
is also a very good conductor of heat, the two properties being 
closely parallel in most cases, a fact which is readily explained 
by the electron theory. 

Silver is too soft for satisfactory use in the free state and is 
usually alloyed with copper to harden it for the manufacture of 
coins or of table ware. “‘ Sterling ”’ silver, used in British coins, 
contains 73 per cent of copper. The coins of the United States 
contain 10 per cent of copper. Silver does not tarnish in moist 
or dry air, but it is easily blackened by hydrogen sulfide or by a 
solution of a sulfide. 

Silver dissolves easily in nitric acid or in hot concentrated 
sulfuric acid, with evolution of nitric oxide, INO, in the former 
ease and of sulfur dioxide in the latter. It is not attacked by 
hydrochloric acid, partly because of the insolubility of the 
chloride. 

Silver Plating. When silver is deposited on a cathode from 
a solution of silver nitrate, or of some other silver salt which is 
ionized largely in solution, it assumes a crystalline form and 
gives a surface which is not suitable for plated ware. If a solu- 
tion of potassium silver cyanide, KAg(CN)s, is used in place of 
one of the ordinary salts, a smooth, coherent deposit can be 
obtained. In the electrolysis the object to be plated is made the 
cathode while a plate of pure silver is used as an anode. The 
object to be plated requires very careful cleaning, as the slightest 
film of grease will cause the deposited silver to flake off. 

Silver Oxide, Ag.O, is formed when a solution of sodium 
hydroxide is added to a solution of silver nitrate. It may be 
supposed that silver hydroxide, AgOH, is formed at first, but 
that this immediately decomposes into silver oxide and water. 
The decomposition seems to represent an equilibrium, however, 
which lies far on the side toward the formation of the oxide, 


SILVER 443 


since the oxide dissolves slightly in water and the solution has 
an alkaline reaction. Also, silver oxide cannot be completely 
freed from water below the temperature at which it dissociates 
rapidly into silver and oxygen. The solution of silver hydroxide 
is ionized to the extent of 70 per cent, indicating that the com- 
pound is a comparatively strong base. 

Silver oxide is easily decomposed by heat, but this reaction, 


also, is reversible : 2 Ap.O%S 4 Ag + O; 


As in all cases of an equilibrium between a gaseous and a 
solid phase, the equilibrium depends only on the temperature 
and the pressure of the gaseous component and not upon the 
relative amounts of the two phases, because the reaction can 
occur only at the surface between the gas and the solid, and not 
throughout the mass of either phase. At 302° the dissociation 
pressure of silver oxide is 20.5 atmospheres. If the pressure is 
decreased, more of the oxide will decompose until all is decom- 
posed or the pressure rises to 20.5 atmospheres again. If the 
pressure is increased, oxygen will combine with the silver till 
all is converted into the oxide or till the pressure falls to 20.5 
atmospheres. At 325° the dissociation pressure is 32 atmos- 
pheres; at 445° it is 207 atmospheres. The dissociation pres- 
sure of silver oxide is 0.2 of an atmosphere at 121°. As 0.2 of 
an atmosphere is the partial pressure of oxygen in the air, it 
follows that at temperatures above 121°, silver oxide will de- 
compose completely if exposed to the air, while at temperatures 
below that metallic silver would change to silver oxide. At 
that temperature, however, both the oxidation and its decom- 
position occur very slowly indeed, if no catalyzer is present 
(Lewis, J. Am. Ch. Soc. 28, 139 (1906)). 

Molten silver absorbs oxygen, which it gives out again, in 
part, as the metal solidifies. It seems probable that the absorbed 
oxygen is not in chemical combination with the silver. 

* Silver Peroxide, AgsOz, is formed as a brown or black coat- 
ing by the action of ozone on silver. It is also deposited on the 
anode as a black, crystalline compound in the electrolysis of 


444 A TEXTBOOK OF CHEMISTRY 


silver nitrate with an electrical pressure of about 15 volts and 
the use of a diaphragm of porous porcelain. 

Silver Nitrate, AgNO3. This is easily prepared by dissolving 
silver in nitric acid and evaporating the solution till the salt 
will crystallize on cooling. It forms tabular crystals of the 
rhombic system. It is easily soluble in water and is used as 
the starting point for the preparation of most of the other com- 
pounds of silver. It melts at 208.6° and is sometimes cast in 
small sticks for use as a cauterizing and germicidal agent in 
medicine. In connection with this use an old name, lunar 
caustic, is still employed — a name which comes to us from the 
alchemists, who recognized a symbolical relation between silver 
and the moon (Latin, luna). 

* Silver Nitrite, AgNOs, is a white, difficultly soluble salt 
easily prepared by precipitating a solution of silver nitrate 
with a moderately concentrated solution of sodium nitrite. It 
is used in water analyses for the preparation of standard solu- 
tions of nitrites. 

* Silver Sulfate, AgoSOu., is prepared by warming silver with 
concentrated sulfuric acid. The salt is only moderately soluble 
(1: 200 in cold water). It is sometimes used to remove chlorine, 
bromine or iodine from solutions when the presence of a nitrate 
is to be avoided. 

Silver Chloride, AgCl, Silver Bromide, AgBr, and Silver 
Iodide, AgI, are almost insoluble compounds which separate as 
curdy precipitates on adding one of the halogen acids or a halide 
to asolution of silver nitrate. The chloride is white ; the bromide, 
yellowish white; and the iodide, yellow. All three of the salts 
are sensitive to light and lose chlorine, bromine or iodine, 
slowly in diffused light, rapidly in sunlight or in artificial light 
which contains rays of short wave lengths — those of the violet 
end of the spectrum or beyond. The action seems to be closely 
related to the effect of sunlight or the magnesium light in causing 
the combination of hydrogen and chlorine (p. 105). 

Photography. Photographic “dry” plates are prepared by 
spreading a thin film of an emulsion of silver bromide in a solu- 


PHOTOGRAPHY. GOLD 445 


tion of gelatin over a plate of glass or of transparent celluloid 
and allowing the film to dry in a dark room. When such a film 
has been exposed to the image formed by the lens of a camera 
no change in the appearance of the film can be noticed, but if 
the exposed plate is placed in a solution of hydroquinone, pyro- 
gallol or some other reducing agent used as a “ developer,” 
those portions of the silver bromide which were struck by the 
light will be reduced to metallic silver, which appears dark and 
opaque, while the portions under the dark parts of the image of 
the camera will not be affected. When the lights and shadows 
have been brought out sufficiently by the developer, the plate is 
placed in a solution of sodium thiosulfate, Na2S.O3, which dis- 
_ solves silver bromide readily, while the metallic silver is not 
affected. If the unchanged silver bromide were not removed, 
it would darken later on exposure to light, and the picture would 
be destroyed. The process of removing the unchanged silver 
bromide is called “ fixing’ the picture. The picture obtained 
is dark where the object photographed was light and light 
where the object was dark. For this reason it is called a “ nega- 
tive.” 

In order to obtain a “ positive,” which will reproduce the 
lights and shadows of the object, the negative is placed over a 
paper which has been coated with a film of silver chloride or 
bromide and is exposed to the light for “ printing.” The light 
passing through the light portions of the negative darkens the 
silver salt beneath, while other portions of the salt are protected 
by the opaque portions of the negative. The picture on the 
paper must be fixed by means of sodium thiosulfate as is done 
with the negative. It may also be “ toned ” by immersion in a 
solution of gold trichloride, AuCls, or of chloroplatinic acid, 
H2PtCle, which will cause the silver to be replaced by gold or 
platinum. ‘The gold gives a reddish tone, the platinum, a steel- 
gray color. 

Gold, Au., 197.2, is found almost always in the free state in 
nature, very rarely in large nuggets weighing several pounds, 
usually in small grains mixed with sand or gravel or dissemi- 


’ 


446 A TEXTBOOK OF CHEMISTRY 


nated through quartzite, granite, pyrite and other rocks and 
minerals. In Colorado and in some other localities it is some- 
times found combined with or alloyed with tellurium in a 
mineral having approximately the composition AuTe:. Al- 
though gold is almost always found mixed with very large 
quantities of worthless or comparatively worthless minerals, it 
is very widely diffused and a careful examination reveals traces 
of gold in almost any rock, soil, clay or other natural, inorganic 
mixture which is tested. It is claimed that sea water contains 
from 0.03 to 0.06 gram of gold in a ton (Liversidge, see Chem. 
Centralblatt, 1905, II, 648). This would be worth from two 
to four cents, but the total amount of gold in the ocean is, of 
course, very large. 

Metallurgy. Gold has a specific gravity of 19.26, while 
that of ordinary minerals averages about 2.6. Such minerals 
may, therefore, be separated from the gold by “ washing ” in a 
current of water, which carries the lighter minerals away, leaving 
the gold. For rich sands or gravels the process may sometimes 
be carried out by hand in a pan — an operation which has given 
the well-known expression ‘‘ pan out.’ On a large scale, in 
“hydraulic mining,” masses of sand and gravel are washed 
away by powerful streams of water, the material running 
through sluiceways in which are placed crosspieces to retain 
the gold. Metallic mercury is usually placed back of the cross- 
pieces, or rifles. This amalgamates with the gold and retains 
it. Dredging is also used to get gold-bearing material into 
sluices. 

Cyanide Process. A few years ago large quantities of gold 
were obtained by a “ chlorination process,” in which the gold 
was dissolved by chlorine obtained from bleaching powder and 
sulfuric acid. This process has been almost completely replaced 
by the cyanide process. The cyanide process has also displaced 
smelting processes in some cases. 

In the presence of air to furnish oxygen, gold dissolves in a 
solution of potassium cyanide as potassium aurous cyanide, 
KAu(CN)>s. Metallic silver and silver chloride will also dissolve. 


GOLD 447 


4Au+8KCN + 0, + 2H.0 = 4KAu(CN),. + 4 KOH 
4Ag+8KCN + 0, + 2H,0O = 4 KAg(CN), + 4 KOH 


The necessity of air in the solution can be shown by shaking 
gold leaf with a solution of potassium cyanide which is free from 
oxygen or through which a current of hydrogen is passed. It 
will not dissolve, while it dissolves readily on passing a current 
of air. Potassium permanganate, potassium chromate, sodium 
peroxide, nitrobenzene or some other oxidizing agent is sometimes 
used to assist in the solution of the gold. The gold is precipi- 
tated from the solution by means of zinc: 


2 KAu(CN)e + Zn = KeZn(CN), + 2 Au 


The solution containing potassium zine cyanide may be used 
for a new lot of ore. 

Native gold frequently contains silver, and gold is often 
obtained along with silver from silver or copper ores. From 
such alloys the silver can be removed by solution in dilute nitric 
acid or in concentrated sulfuric acid, provided not more than 
35 per cent of gold is present. Alloys which are richer in gold 
than this are melted with enough silver to reduce the amount of 
gold to one third or one fourth of the whole. The separation 
with concentrated sulfuric acid can be carried out in cast-iron 
kettles and is cheaper than that with nitric acid. It also has 
the advantage that any platinum which Js present remains with 
the gold, while small amounts of platinum dissolve with the silver 
in nitric acid. The separation of gold from silver by solution of 
the silver is called, technically, “ parting.” 

The production of gold in the United States in 1910 was 
4,657,018 troy ounces worth $96,269,100. The coining value 
of an ounce of gold is $20.67183. The average yearly production 
of gold in the world during the first half of the nineteenth cen- 
tury was only about 800,000 troy ounces. The present annual 
production in the United States alone is nearly six times that 
and the annual production in the world is many times that of 
70 or 80 years ago. 


448 A TEXTBOOK OF CHEMISTRY 


Properties of Gold. Gold is yellow by reflected light but 
transmits the complementary color, green, through very thin 
films, as through gold leaf held between two glass plates. It is 
the most ductile and malleable of all the metals and can be 
drawn into exceedingly fine wires and beaten into very thin 
leaves. Its specific gravity is about 19.3, varying considerably 
with the method of preparation and treatment. It melts at 
1063°. Its electrical conductivity is two thirds that of silver. 

Gold is insoluble in any one of the common acids, alone, but 
it dissolves readily in aqua regia, a mixture of three volumes of 
hydrochloric acid with one of nitric. Its solubility in chlorine 
water has been mentioned above. It dissolves also in selenic 
acid. 

Alloys of Gold. Gold is a very soft metal and is alloyed with 
other metals, especially with silver and copper for use in jewelry 
and coins. British gold coins are 22 carats fine, 2.e. #4 pure 
gold. In reference to gold the term “ carat ”’ is used to designate 
the number of parts in a total of 24 which consist of pure gold. 
Thus 18-carat gold is 3% or ? pure gold. American gold 
coins are made on a decimal basis ‘‘ 900-fine,”’ 7.e. 900 parts in 
1000 are pure gold. The American eagle contains 900 parts of 
gold and 100 parts of copper. 

* Oxides of Gold. Three oxides of gold have been described : 
gold monoxide, AuszO; gold dioxide, AuO; and gold trioxide, 
Au2,O3. There seems to be some doubt whether the second of 
these has been prepared as a definite compound. Each oxide is 
decomposed at a comparatively low temperature into gold and 


oxygen. 


ZO 
* Gold Hydroxide, ANZO2E, may be prepared by adding 
magnesium carbonate, MgCOs, to a solution of chloroauric acid 
and dissolving the excess with dilute nitric acid : 


HAuCl, + 2 MgCO; = HAuO, + 2 MgCl, + 2 CO, 


Gold hydroxide is a yellowish brown precipitate having the 
properties of both a base and an acid. It dissolves in hydro- 


ELECTRICAL CONDUCTANCES 


449 


Speciric ConDUCTANCE AND RESISTANCE OF Common METALS! 











Aluminium 
Bismuth 
Cadmium 
Cobalt 
Copper 
Gold Fay 
Iron (electrolytic) 
Iron (steel, 1% C) 
Lead 
Magnesium . 
Mercury . 
Molybdenum 
Nickel 
Palladium 
Platinum . 
Potassium 
Silicon 
Silver . 
Sodium 
el in eerese ae Tt 
Tungsten  (fila- 
ment) 
Zine 


TEMPER- 


RESISTANCE IN 

















SPECIFIC SPECIFIC OMS OF WIRE 
ATURE CONDUCTANCE RESISTANCE 1m. Lone 1 mM. 
IN DIAMETER 
eee o0.O 7 Se lO 2.8 he cel ae 0.036 
18° ONS 5-6 Mion Sevidiesy sealt bs 1.58 
ee Ler ex erb.8o LOT” 0.087 
BOSE PIO .S 8 SOLON L708 0.123 
257 GO V8. Gao 107 21271 X1105% 0.022 
Ose 15 CLO TO On One 0.027 
a Steere Os aL La xed Ore 0.154 
18° BU ama er LO un 10,0 xpd One 0.254 
0° BAe 10% 10:4 ese 07% 0.247 
Q2e eres ex a sco 1 (© 0.055 
0° 1.063 <10°7.94.07 «107° E192 
2 caeaal 00 ee Oo 5:6 ox LOR® 0.071 
ae gah te 1051 9. 1e cul One 0.116 
0° Cala UA 1U.6 axel) @ 0.135 
20 ee Ue Oe ee Oo Le 0.118 
Peeee LA eed Oe ban G, (2 oO 0.086 
— POs LUCID OD joctlOn: 0.76 
ape WARES pret ileal Ma ayal atl Eee 0.020 
Terme oO Spo 1 Os Wadia x<e10 ao 0.061 
18° See 10h 18 10% 0.169 
COTE mee 1OFSD 6 BOS 10 0.071 
1G Gay LG Oreo OF 86.0 810 %* 0.076 

















1 Specific conductance is the current in amperes which would 
pass between opposite faces of a cubic centimeter of the metal under 


a potential difference of one volt. 


reciprocal of the specific conductance. 
the specific resistance is given in microhms, 7.e. in millionths of 
anohm. The resistance of a wire 1 m. long and 1 mm. in diameter 


is found by multiplying the specific resistance by 10000 x 


The specific resistance is the 
If the factor 10~* is omitted, 


3.1416 


450 A TEXTBOOK OF CHEMISTRY 


chloric acid with regeneration of chloroauric acid. It also dis- 
solves in a solution of potassium hydroxide and from the solution 
potassium aurate, KAuQ:.3 H,O, may be crystallized. 

Chlorides of Gold. Corresponding to the three oxides there 
are three chlorides of gold: gold monochloride, AuCl; gold 
dichloride, AuCl.; and gold trichloride, AuCl;. The last of these 
combines with hydrochloric acid to form chloroauric acid, 
HAuCly, and this last compound is formed when gold is dissolved 
in aqua regia. It is deposited by evaporating and cooling such. 
a solution, in the form of yellow needles having the composition 
HAuCl..4 H.O. It is a monobasic acid, which forms well- 
defined, crystalline salts with many of the metals and especially, 
also, with many organic bases. The potassium = salt, 
KAuCl.2 H.O, the calcium salt, Ca(AuCl,)s.6 H.O, the 
ammonium salt, 2 NH,AuCl4.5 H2O, and the strychnine salt, 
Co:H22N202.HAuClh, are typical of such compounds. 


EXERCISES 


1. Write the equation representing the action of nitric acid on copper ; 
also for the action of nitric acid on cupric sulfide. Free sulfur and nitric 
oxide are formed in the latter case. 

2. Write the equations for the action of nitric acid and of sulfuric 
acid on silver. 

3. How many cubic centimeters of nitric acid of specific gravity 1.42, 
containing 70 per cent of the pure acid, will be required to dissolve a 
pound (453 grams) of silver ? 

4. How many grams of concentrated sulfuric acid of 98 per cent will 
be required to dissolve a pound of silver, taking account only of that 
which enters into the reaction ? 

5. Write the equation for the action of potassium iodide on copper 
sulfate; also for the reaction between sodium thiosulfate and iodine. 
How many milligrams of crystallized sodium thiosulfate, Na2S.O03.5 H.O, 
are equivalent to 63.6 milligrams of copper in these reactions? How 
many milligrams are equivalent to one milligram of copper? 

6. Write the equation representing the solution of gold in selenic 
acid, H»SeO,. The gold dissolves in the trivalent form, while some of 
the acid is reduced to selenious acid. 


CHAPTER XXVI 


GROUP II, ALKALI-EARTH METALS: BERYLLIUM, CALCIUM, 
STRONTIUM, BARIUM, RADIUM 


THE elements calcium, strontium and barium are called alkali- 
earth metals, a somewhat indefinite designation coming down to 
us from the time of the alchemists and referring to the fact that 
the hydroxides of these metals are strong bases. The elements 
beryllium and radium are not included under the designation. 
All of the elements of Group II are bivalent. Calcium, stron- 
tium, and barium decompose water at ordinary temperature, 
though much less rapidly than the alkali metals. Calcium 
hydroxide is difficultly soluble in water. Strontium and barium 
hydroxides are somewhat more soluble. All three carbonates 
are nearly insoluble in water, but dissolve as bicarbonates in 
water containing carbon dioxide. Barium sulfate is one of the 
most insoluble salts, while calcium sulfate is slightly soluble 
(1 part in 500 of water). 

* Beryllium,! Be, 9.01. The mineral bery] is a silicate of beryl- 
lium and aluminium, having the composition Be3Al:SigQig (or 
3 BeO.Al2.03.6 SiOz). As a precious stone, different forms of 
the mineral are known as emerald and aquamarine. The free 
element is silver-white and has a specific gravity of 1.9. Both 
as an element and in its compounds beryllium resembles mag- 
nesium and aluminium rather than calcium. The hydroxide, 
Be(OH)., is nearly insoluble in water but dissolves both in acids 
and in alkalies, exhibiting in this way both acid and basic 
properties. The chloride, BeCl,.4 H2O, sulfate, BeSO..4 HO, 


1 Sometimes called glucinum, Gl, a name which has some claim 
to priority. The name beryllium has the advantage of referring to 
the most important mineral containing the element. 


451 


452 A TEXTBOOK OF CHEMISTRY 


and nitrate, Be(NO3)e.3 H2O, are soluble in water. The car- 
bonate is unstable and readily loses carbon dioxide, forming a 
Hasic carbonate. 

Calcium, Ca, 40.07. Occurrence. Calcium is one of the most 
abundant and most important of the elements. It is found as 
the carbonate, CaCQOs, in the minerals calcite and aragonite, in 
marble and, less pure, in limestones ; as the sulfate in the mineral 
gypsum, CaSO,.2 H,O; as the phosphate, Ca3(POx)2, in bone 
ash and mineral phosphates and in the mineral apatite, 
Ca;(PO,4)3F ; and as the fluoride in the mineral fluorite, CaF». 

Preparation, Properties. Metallic calcium can be prepared 
most easily by the electrolysis of the fused chloride (E. F. Smith 
and Goodwin, J. Am. Chem. Soc. 25, 873 (1903)). It is a white, 
crystalline metal which decomposes water readily at ordinary 
temperatures, with the formation of the hydroxide, Ca(OH)s. 
It burns in air, forming both the oxide, CaO, and the nitride, 
CazNeo. The latter is hydrolyzed by water with the formation of 
ammonia : 

CasN2 + 6 HOH = 3 Ca(OH): + 2 NH3 

Calcium melts at 800°. 

* Calcium Hydride, CaHg, is a white, crystalline compound 
which can be prepared by the direct union of the elements. 

Calcium Oxide, CaO, or Lime is manufactured on a large 
scale by heating the carbonate in lime kilns. In the older forms 
a mixture of limestone and fuel was placed in the kiln. The 
fuel was then set on fire and allowed to burn till it was all con- 
sumed. In the newer forms a mixture of limestone and coal is 
charged into the kiln at the top, while the “ burnt” lime is 
removed at the bottom, without stopping the process. The 
dissociation of the carbonate: 


CaCO; = CaO + COs, 


is a reversible reaction. As there is a gaseous constituent, the 
system has a characteristic pressure for each temperature. 
This may be determined by heating calcium carbonate in a plat- 
inum bulb connected with a manometer. The pressure of the 


ALKALI-EARTH METALS: CALCIUM 453 


carbon dioxide which will be in equilibrium with a mixture of 
calcium oxide, CaO, and calcium carbonate is as follows: 


Temperature 600° 700° 800° 898° 950° 
Pressure of CO, 2.35 mm. 25.3 mm..168 mm. 760 mm. 1490 mm. 


It will be seen from the table that the temperature of the lime 
kiln must be above 900° for the carbon dioxide to escape rapidly 
from the interior of a piece of limestone. When calcium car- 
bonate is heated in an open crucible in the laboratory under 
conditions such that the carbon dioxide is constantly displaced 
by air, the partial pressure of the carbon dioxide may become very 
low and complete decomposition could be secured at a tempera- 
ture of 750° or below. 

In the transformation to lime, pieces of limestone retain their 
shape but shrink somewhat in size. If water is added to the 
lime, it combines with it, evolving a very considerable amount of 
heat — 15,540 small calories per gram molecule. At the same 
time the calcium hydroxide, Ca(OH)2, which is formed, swells 
and falls to a loose powder. The process is called “slaking.” If 
lime is exposed to the air, it slowly absorbs water and falls to a 
powder, and the hydroxide also absorbs carbon dioxide and is 
converted back to the carbonate. Because of the latter change, 
“air-slaked lime”’ is usually worthless for the preparation of 
mortar. 

* Dissociation of Calcium Carbonate and the Phase Rule. 
In the system produced by the partial dissociation of calcium 
carbonate there are three phases (CaCQO3, CaO and COs, and only 
two components (CaO and COz), the calcium carbonate formed 
by the union of the other two not being considered as a separate 
component. According to the phase rule (p. 107) a system hav- 
ing two components and three phases has only one degree of free- 
dom and is univariant, just as the system water vapor — water 
has one component and two phases and is also univariant. 
Accordingly, in the system calcium carbonate — calcium oxide 
—carbon dioxide, if the temperature changes, the pressure 
must change also; and for every temperature there is a fixed 


454. A TEXTBOOK OF CHEMISTRY 


dissociation pressure, just as for every temperature there is a 
fixed vapor pressure for the system water vapor — water. 

Mortar is prepared by mixing slaked lime with water and sharp 
sand, which has not been rounded by long action of waves. 
Where used between bricks the moisture is partly absorbed by 
the bricks and partly dries out in the air. This is followed by 
the action of the carbon dioxide of the air, which slowly changes 
the hydroxide to calcium carbonate: 

Ca(OH). + CO, == CaCO; + H,.O 

The carbonate crystallizes as it forms and adheres strongly to 
the particles of sand, binding them together. In plastered 
rooms the liberation of moisture by the reaction keeps the air 
of the rooms moist for some days or weeks. 

Cement. For the manufacture of cement a clayey limestone 
is sometimes used, but, as limestones having the proper composi- 
tion are rare, artificial mixtures of finely powdered limestone and 
a fine clay or shale, rich in silica, are usually employed. The 
slag from blast furnaces is also extensively used. The materials 
are finely ground and fed into the top of a long, slanting, slowly 
rotating, tubular furnace which is heated by means of powdered 
coal blown in at the lower end. A temperature of 1500°-1600° 
is reached in the hottest parts of the furnace. The carbon 
dioxide is completely expelled and a “ clinker’”’ composed of 
calcium silicate and calcium aluminate with an excess of calcium 
oxide is formed. This clinker is finely ground and mixed with 
a small amount of plaster of Paris. It should contain only a 
very small per cent of magnesium. The composition of the 
finished cement is as follows: 


Loss on ignition 0-2 per cent 
Silica, SiO» 15-20 per cent 
Alumina, Al,O3 3-8 per cent 
Ferric oxide, Fe.O; 3-6 per cent 
Lime, CaO 58-64 per cent 
Magnesia, MgO 0-4 per cent 
Potash and soda, K,O, Na2,O Q-2 per cent 


Sulfur trioxide, SO; 0-2 per cent 


CALCIUM CHLORIDE 455 


When the cement is mixed with water it slowly combines with 
it, forming partly crystals of calcium hydroxide, Ca(OH)s, partly, 
probably, hydrated silicates and aluminates of calcium which 
set”? to a firm, strong mass. Sand and sometimes other 
materials are usually added to increase the volume. As the set- 


ce 


ting results from the action of water alone, it will take place under 
water and hence the material is often called hydraulic cement. 

Calcium Chloride, CaCl2, may be prepared by dissolving cal- 
cium carbonate in hydrochloric acid, evaporating the solution 
and drying the residue, finally at a temperature of 260° or above. 
Prepared in this manner it forms a porous, extremely hygro- 
scopic and deliquescent mass, which is much used for drying 
gases. The anhydrous salt melts at about 800°. With water 
it forms a series of hydrates. The one which is stable at ordinary 
temperatures and which may be crystallized by cooling very 
concentrated solutions is CaCly.6 H,0. Considerable heat is 
evolved when the anhydrous chloride dissolves in water, but a 
mixture of the hydrate with a little less than its weight of snow 
will give a lowering of the temperature to — 55°. 

Even the hydrate, CaCly.H.2O, has an appreciable vapor pres- 
sure at ordinary temperatures, and gases cannot be fully dried 
by the use of the anhydrous salt. At 25° it leaves 0.36 
milligrams of water in liter of a gas which has been passed 
‘over it. 

Calcium chloride is obtained in large quantities as a by-product 
in the ammonia-soda process (p. 412). Because of the low 
freezing point, solutions of calcium chloride are used in refrigerat- 
ing machines to surround the cans in which water is frozen to 
artificial ice. They have also been used for sprinkling roads. 


SOOl 


“ Chloride of Lime,” OES , ls prepared commercially by 
Cl 


‘spreading slaked lime on the floor of a room and filling the room 
with chlorine gas: 


OCl 
eal 


Ca(OH). a Cl, => BE -- HO 
‘el 


456 A TEXTBOOK OF CHEMISTRY 


Treatment with a dilute acid liberates hydrochloric and hypo- 
chlorous acids and these, in turn, react to form free chlorine 
(p. 126). The compound is used for bleaching (p. 125) and for 
sterilizing water (p. 83). Good bleaching powder should 
contain 35-37 per cent of “‘ available” chlorine. It deteriorates, 
partly by loss of hypochlorous acid through the action of the 
carbon dioxide of the air, partly by a slow transformation of 
the hypochlorite to the chlorate. 

* Calcium Chlorate, Ca(ClO3)e, is prepared by passing chlorine 
into “ milk of lime,” a mixture of calcium hydroxide and water. 
The hypochlorite formed at first changes to chlorate by auto- 
oxidation. The solution is used to mix with a solution of potas- 
sium chloride, KCl, for the preparation of potassium chlorate, 
KCIO3. (Why ?) 

Calcium Fluoride, CaF, is found as the mineral fluorite, which 
crystallizes in cubes or octahedra. It is nearly insoluble in water, 
and for this reason natural waters never contain more than 
minute traces of fluorine. Calcium fluoride melts at 1330°. 
It is used as a flux in casting iron and in other metallurgical opera- 
tions and is the source from which hydrofluoric acid and all other 
compounds of fluorine are obtained. 

Calcium Sulfide, CaS, is formed by the reduction of calcium 
sulfate, CaSOx., by heating it with coal or charcoal in the Leblanc 
Soda Process (p. 411). It does not seem to be appreciably 
soluble in water, but is slowly hydrolyzed to the hydrosulfide, 
Ca(SH)o, which dissolves : 


2 CaS + 2 HOH = Ca(SH). + Ca(OH), 


The hydrolysis is sufficiently slow in an alkaline solution so 
that the sodium carbonate of the Leblane process may be dis- 
solved, leaving nearly all of the calcium sulfide as an insoluble 
residue. 

In the earlier manufacture the calcium sulfide was discarded 
as a waste product, but the slow hydrolysis of the material on 
exposure to the weather led to the contamination of streams in 
England and resulted in stringent legislation requiring the 


PLASTER OF PARIS 457 


manufacturers to take care of their waste products in such a 
manner that they should not become a nuisance to others. 
This and the competition of the ammonia-soda process, com- 
pelling the manufacturer to practice every possible economy, 
led to the invention of the Chance process for the recovery of 
the sulfur. By exposing the moist calcium sulfide to the action 
of carbon dioxide, the calcium is converted to the carbonate and 
hydrogen sulfide is liberated : 


‘Ca(SH)2 + H2CO; = CaCO; + 2 HS 


The fact that the ionization constant of carbonic acid is con- 
siderably greater than that of hydrogen sulfide and also the 
insolubility of the calcium carbonate, both aid in shifting the 
equilibrium of the reaction to the right. 

By burning the hydrogen sulfide with a limited supply of air 
it is possible to convert the sulfur almost completely to the free 
state. 

* Acid Calcium Sulfite, CaH2(SOs3)s, is prepared by burning 
sulfur and passing the sulfur dioxide, SO2, formed into milk of 
lime. The solution is used to dissolve and remove lignin from 
the fiber in the manufacture of paper from wood. 

Calcium Sulfate, CaSO,.2 H.O, Plaster of Paris. The min- 
eral gypsum, which has the composition CaSQ.4.2 H.O, is found 
in nature in clear, transparent, monoclinic crystals called selenite 
and in a white, opaque form called alabaster, also in large quan- 
tities in a form suitable for making plaster of Paris. Calcium 
sulfate is also sometimes found in nature in the anhydrous form 
in the mineral anhydrite. If gypsum is treated for some time 
at 130°-160°, it loses three fourths of the water which it contains 
and is converted into a compound having the composition 
2 CaSOy.H.O and known commercially as plaster of Paris. 
When this is mixed with a small amount of water to a creamy 
consistency, itcan be filled into a mold; but after standing a short 
time part of the water unites chemically with the salt to form 
gypsum, CaSQ,.2 H.O, which “ sets”’ to a solid mass. If the 
plaster is heated to too high a temperature or for too long a 


458 A TEXTBOOK OF CHEMISTRY 


time in driving out the water, it becomes “‘ dead burnt ” and will 
then combine with water only very slowly, and it is worthless 
for the ordinary uses of plaster of Paris. It seems probable that 
ordinary plaster of Paris retains a small amount of the unchanged 
dihydrate and that the molecules of this furnish the starting 
point for the crystallization in setting. Calcium sulfate is 
difficultly soluble, about two grams dissolving in a liter of water. 
It is much less soluble in alcohol. 

* Plaster of Paris and the Phase Rule. In the preceding 
paragraph two hydrates of calcium sulfate, CaSO4.2 H.O and 
2 CaSO,.H2O, have been mentioned and also a natural anhy- 
dride, the mineral anhydrite. Another anhydride, which is 
more easily soluble than the natural anhydrite, is also known and 
is called the “soluble anhydrite.” In speaking of the vapor 
pressure of hydrates (p. 82) it has been implied that each 
hydrate has a characteristic vapor pressure. This is strictly 
true only in case the loss of water leads to the formation of only 
one compound, either a lower hydrate or an anhydride. The 
majority of salts lose water in this manner; but gypsum, 
CaSO,.2 HO, may lose water in such a manner as to form either 
of the three substances, natural anhydrite, soluble anhydrite or 
plaster of Paris, 2 CaSO..H,O. The vapor pressure of the 
gypsum will depend on which of the three substances is formed, 
as is seen in the table on the opposite page. 

If these values are plotted and the vapor pressure curves 
extended, it is found (Fig. 101) that the curve for the system 
CaSOx..2 H,O, CaSO, (natural anhydrite) cuts the curve for 
water vapor at 66° and the other two curves cut it at 89° and 
107°. As there are only two components (H2O and CaSQs,), 
each of these temperatures represents a quadruple point (p. 78) 
where the sysem is invariant. The four phases at 66° are: 
vapor, solution, CaSO..2 H:O and CaSO, (natural anhydrite). 
Any change in temperature or in pressure will cause the disap- 
pearance of one of the phases. If the pressure is decreased, 
water will evaporate till only the natural anhydrite is left. If 
the pressure is increased, the vapor phase will disappear. If the 


CALCIUM SULFATE: PHASE RULE 459 


Vapor PRESSURE IN MILLIMETERS or Mercury FoR SYSTEMS 
CoNTAINING GypsuM 

















SysTpM SysTEmM SYSTEM 
Sen ie | eee asm Peswige 
ANHYDRITE ANHYDRITE 2 CaSO. HO 

15 ies 8.43 rf 4.2] 
20 17.4 122 10.7 6.24 
3 31.5 24 19.4 2a 
40 54.9 45.4 o4 26.3 
50 149 143 108 91.4 
65 187 = 140 122 
70 233 —— 185 161 
SO 300 —— 314 272 
90 526 — SSS 446 

109 760 —— — 711 

105 906 —— SS 888 

MLO MRELO Lo = ae a = 







































































460 A TEXTBOOK OF CHEMISTRY 


temperature is increased, all of the gypsum, CaSO..2 HO, 
will be changed to anhydrite because the vapor pressure of 
gypsum would be greater than that of anhydrite above this 
temperature. The change might, it is true,.take a long time, but 
there would be no stable equilibrium till the change was com- 
plete. If the temperature was lowered, all of the anhydrite 
would be converted into gypsum, since the vapor pressure of 
water is greater than that of gypsum at temperatures below 66°. 

A consideration of the curves will enable one to predict whether 
gypsum or an anhydrite will crystallize from a salt solution, if 
we know the vapor pressure of the latter. Thus a solution of 
sodium chloride has a vapor pressure greater than 12.2 mm. at 
20°. From such a solution gypsum will crystallize. If mag- 
nesium chloride or calcium chloride is added till the vapor pres- 
sure is less than 12.2 mm. at 20°, anhydrite will crystallize from 
the solution, because if gypsum and such a solution were placed 
side by side in a confined space water would escape from the 
gypsum and condense in the solution. These predictions of the: 
theory agree with facts observed by geologists about the condi- 
tions under which gypsum and anhydrite are found in nature. 

* Calcium Nitrate, Ca(NOs)2.4 H2O, is readily prepared by 
dissolving the carbonate in nitric acid. It is so manufactured 
for fertilizers and other uses on a considerable scale in Norway, 
by absorbing the oxides of nitrogen, formed by the electric arc 
from air, in milk of lime. The anhydrous salt is sometimes used 
to dry oxides of nitrogen or other gases for which calcium chloride 
cannot well be used. 

Calcium Phosphates. Normal calcium phosphate, Ca3(POx)2, 
is an important constituent of bones. It is also found as a min- 
eral phosphate in deposits in the southeastern part of the 
United States from North Carolina and Tennessee to Florida. 
Both bones and the mineral phosphates are extensively used as 
fertilizers, phosphorus being an element which is essential for 
the growth of crops and which is found in only limited amounts 
in some soils. In order to render the phosphate more easily 
soluble and available for the growth of plants, the powdered 


CALCIUM PHOSPHATES 461 


mineral phosphate is often treated with sulfuric acid to convert 
it into monocalcium phosphate, CaH4(POx,)o. The mixture of 
calcium sulfate and acid calcium phosphate is designated, com- 
mercially, as a “superphosphate.” A slightly diluted acid is 
used and this “‘superphosphate ”’ contains both salts as hydrates : 


Ca3(PO,)2 + 2 H.SO, oe 6 H.O 
— Ca(H2PQ,)>2.2 H.O + 2 (CaSOx4.2 H,.0) 


* Solubility of Calcium Phosphates. Even the monocalcium 
phosphate, Ca(H2PQO.)s, is only slightly soluble, about 5 grams 
dissolving in a liter of water. If a larger amount of the salt is 
added, the reaction : 


2 H.POs = HPOs— + H3PO, 


causes the number of monohydrophosphate ions, HPO, _, and 
of calcium ions in the solution to exceed the solubility product for 


the reaction, _ Cat+ + HPO, ~ = CaHPO, 


and there will be formed a precipitate of the very difficultly 
soluble dicalcium phosphate, CaHPQO,, while the solution will 
contain more phosphoric acid than will correspond to the mono- 
calcium phosphate, Ca(H2POx.)2. In the valuation of fertilizers 
it is customary to distinguish three forms of phosphoric acid, 
“water-soluble phosphoric acid,” “‘ citrate-soluble phosphoric 
acid ”’ (phosphoric acid insoluble in water but soluble in a neutral 
solution of ammonium citrate of sp. gr. 1.09) and “ insoluble 
phosphoric acid.” The last is supposed to be in the form of 
tricalcium phosphate, Ca3(PO,)2. From what has been said 
above it is evident that more phosphoric acid will pass into 
solution if a dicalcium phosphate is treated with successive small 
portions of water than if it is treated at once with a large quan- 
tity of water. The amount of the “ citrate soluble phosphoric 
acid ’’ also depends on the exact conditions of the determination. 
Both determinations are to be considered as in a considerable 
measure conventional, and it is quite certain that they do not 
furnish an accurate measure of the availability of the phosphorus 


462 A TEXTBOOK OF CHEMISTRY 


for the growth of plants. Finely ground phosphate rock also 
furnishes phosphorus which can be slowly absorbed by plants, 
and it is at least a question whether the acid phosphate is any 
better than the raw ground phosphate for maintaining the 
fertility of land during a series of years. 

The addition of even a weak acid, such as acetic acid, to cal- 
clum phosphate will cause it to pass into solution. This is 
because phosphoric acid is so weak an acid that in the presence 
of even comparatively few hydrogen ions, H*, the phosphoric 
acid must be almost completely either in the form of the un-ion- 
ized acid, H3POu, or of the dihydrogen phosphate ions, H,PQ,-. 
Under these conditions there can not be enough of the mono- 
hydrogen phosphate ions present to reach the solubility product : 


Caex/HPO,- -— 


which would cause the precipitation of the dicalcium phosphate. 
As has been explained above, it is this precipitation which 
causes the apparent difficult solubility of monocalecium phos- 
phate. 

Calcium Carbide, CaCs. By heating a mixture of lime, CaO, 
and coke in a revolving electric furnace, calcium carbide is 


formed : GaO 8 G=tCaG et CO 


The carbide is easily hydrolyzed by water with the formation 
of calcium hydroxide and acetylene, CH» (p. 293), and is manu- 
factured chiefly for that use. 

Calcium Cyanamide, CaCN»2. By heating calcium carbide 
in a current of nitrogen at 1000° it is transformed into calcium 


cyanamide, Ca—N—C=N: | 
CaC, + 2 N = CaCN, + C 


Calcium cyanamide is hydrolyzed by water to ammonia and 
calcium carbonate : 


CaCN, -}- 3 H.O = CaCO3 a 2 NHs3 


As these reactions furnish a means of transforming the nitrogen 
of the air into a form which is available for plant growth, calcium 


HARD WATERS 463 


cyanamide is now manufactured in considerable quantities for 
use as a fertilizer. For this use it is often called “ lime-nitro- 
gen,” or in German “ Kalk-Stickstoff.” 

Calcium Carbonate, CaCOs, is the most abundant compound 
of calcium. ‘The various forms have been already mentioned, 
also its conduct when heated. 

Hard Waters. One liter of pure water at ordinary tempera- 
tures will dissolve only about 12 milligrams of calcium carbonate, 
CaCOs3, but water saturated with carbon dioxide will dissolve 
nearly 100 times as much, or more than a gram in one liter. 
The solution contains the acid carbonate, Ca(HCOs)., frequently 
called the bicarbonate. As natural waters always contain some 
carbon dioxide absorbed from the air and usually acquire much 
more from decaying vegetable matter in the soil, all such waters 
which have come in contact with a soil containing calcium car- 
bonate hold more or less of the caletum bicarbonate in solution. 

The properties of such a solution can be easily illustrated by 
passing carbon dioxide through a solution of limewater, slightly 
diluted. Calcium carbonate will be precipitated at first, but on 
continuing the current of the gas it will pass again into solution. 
On boiling the solution the acid carbonate dissociates into 
calcium carbonate, carbon dioxide and water. The carbon 
dioxide escapes with the steam and the calcium carbonate is 
precipitated. Natural waters of this type are said to have 
“temporary hardness,’’ since the hardness is nearly all removed 
by boiling. The designation “hardness” refers to the effect 
of hard water in precipitating an insoluble calcium salt when 
soap is used with it. The water continues to have a harsh feel- 
ing to the skin until enough soap has been used to complete the 
precipitation of the calcium and magnesium salts which are 
present. : 

Waters containing calcium sulfate in solution will not deposit 
the salt on short boiling and are said to have “ permanent 
hardness.” Such waters deposit the sulfate on concentration 
of the water, as is done in a steam boiler, and the decreased 
solubility of the calcium sulfate at a high temperature increases 


464 A TEXTBOOK OF CHEMISTRY 


the amount of scale formed. The scale from such a water is 
especially coherent and troublesome. 

If milk of lime, Ca(OH )s, is added to water containing calcium 
bicarbonate in just the right proportion, nearly all of the calcium 
carbonate in the water and also the calcium of the milk of lime 
which is added will be precipitated. This is known as Clark’s 
process of softening water : 


Ca(HCQs3)2 + Ca(OH): = 2 CaCO; + 2 H:O 


To remove the calcium of calcium sulfate, sodium carbonate, 
trisodium phosphate, sodium fluoride or some other salt which 
will precipitate the calcium must be used. 

* Determination of Free and Combined Carbonic Acid in 
Natural Waters. It has been shown (p. 389) that phenol- 
phthalein is a suitable indicator for weak acids; and methyl 
red, for weak bases. If the acid is very weak indeed, the 
hydrolysis of the normal salt may cause the end point with 
phenolphthalein to appear before the acid is fully neutralized. 
In the case of carbonic acid, H2CO3, when the point correspond- 
ing to the formation of NaHCQs is passed, the hydrolysis of the 
sodium carbonate, 


Nat + Nat + CO;- + H* + OH™ 
= Na* + Nat + HCO; + OH- 


causes the concentration of hydroxide ions to exceed the end 
point for phenolphthalein. The amount of alkali required to 
give a pink color with phenolphthalein will, therefore, indicate 
the amount of free carbonic acid, H2COs, in the solution. 

On the other hand, if a strong acid is added to a solution of a 
carbonate or bicarbonate, so long as any of either remains in 
solution the concentration of the hydrogen ions cannot exceed 
the concentration in a solution of carbonic acid, H2CO3. When 
enough acid has been added to decompose all of the carbonates 
and bicarbonates present, any further addition of acid will 
carry the concentration of the hydrogen ions past the end point 
for methyl orange or methyl red. This mekes it possible to 


CALCIUM OXALATE 465 


determine the amount of free carbonic acid and also of carbonates 
and bicarbonates present, by titrating first with alkali, using 
phenolphthalein as an indicator and then with an acid, using 
methyl orange or methyl red. 

* Calcium Acetate, Ca(C2H3O2)s, is prepared commercially by 
neutralizing the distillate obtained by the destructive distilla- 
tion of wood. It is used for the manufacture of glacial acetic 
acid, HC2H3O2, and of acetone, CH3COCHs3. . 

Calcium Oxalate, CaC.O4.H,O. When a solution of ammo- 
nium oxalate, (NH4)eC2Ou, is added to a solution containing a 
soluble salt of calcium, calcium oxalate is precipitated as a fine, 
crystalline powder. One liter of water dissolves only 5.6 
milligrams of the salt, and it is still less soluble in a solution con- 
taining ammonium oxalate. For this reason it is often used for 
the quantitative determination of calcium. Its value for this 
purpose is greatly enhanced by the fact that magnesium oxalate, 
MegC,O,, is much more easily soluble — 300 milligrams in one 
liter of water. 

In strong acids, as hydrochloric or nitric acid, calcium oxalate 
dissolves. Oxalic acid is only moderately ionized in solutions of 
medium, concentration — 50 per cent to Ht and HC.O, 7 in 
tenth normal solution. The ionization to H*, H* and C,0,-~ 
must be very much less. As the presence of the hydrogen ions 
of a highly ionized acid, such as hydrochloric acid, shifts the 
equilibrium of the reaction : 


H.C204 ia H* + Hay ste C.0.— 


to the left, the concentration of the oxalate ions, C204, in such 
a solution cannot be great enough for the solubility product 
C.0,-~ X Ca** to reach the point of precipitation for calcium 
oxalate. The salt will, therefore, dissolve in such a solution. 
In a solution of a weak acid such as acetic acid, HC2H30., the 
number of hydrogen ions is so small that they produce only a 
slight effect on the ionization of oxalic acid, which is a very 
much stronger acid than acetic acid. Calcium oxalate may be 
precipitated, for this reason, from solutions containing acetic 


466 A TEXTBOOK OF CHEMISTRY 


acid, though the salt is more soluble in dilute acetic acid than 
in pure water. Owing to the extremely low ionization constant 


fon HPO. = Ht + HPO, 


calcium phosphate is not precipitated from solutions containing 
even the very weak acid, acetic acid. The addition of ammo- 
nium oxalate to a solution of calcium phosphate in dilute acetic 
acid will, accordingly, cause the precipitation of nearly all of the 
calcium in the form of the oxalate, and even a solution of calcium 
sulfate will precipitate oxalic acid, but not phosphoric acid, from 
a solution of an oxalate containing acetic acid. 

Calcium oxalate loses carbon monoxide and is converted 
into calcium carbonate by gentle ignition: 


CaCO, = CaCO; a CO 


Calcium Silicate, CaSiO3, is found in nature as the mineral 
wollastonite. It may be prepared by fusing a mixture of quartz 
or sand with calcium carbonate : 


CaCO3 + SiO» = CaSiO; + CO 


Calcium silicate is a constituent of a very large proportion of 
the most common natural silicates, such as pyroxene, amphibole, 
garnet and the zeolites. 

Glass. By melting together a mixture of calcium carbonate, 
CaCO3, sodium carbonate, NasCO;, and a pure quartz sand, 
SiO., in proper proportions, a silicate of calcium and sodium is 
obtained, which does not readily crystallize on cooling, but which 
passes through a stage in which it becomes more and more viscous 
and finally solidifies to a transparent, homogeneous mass. ‘The 
sodium may be partly or completely replaced by potassium, the 
calcium may be replaced by lead, and part of the silica by boric 
anhydride, B2O3, giving glasses suitable for special uses. In all 
of these the glass is to be considered as a complex mixture of 
silicates in the form of an extremely viscous, supercooled liquid. 
The value of glass depends largely upon the fact that as an 
amorphous, supercooled liquid it is still viscous but sufficiently 


GLASS. STRONTIUM 467 


plastic so that it can be worked over a considerable range of 
temperature. ; 

Some of the more important varieties of glass are as follows: 
window glass, plate glass and the glass of ordinary table ware are 
usually a silicate of calcium and sodium. ‘The finer grades of 
such glass are often called crown glass. Flint glass is a silicate 
of lead and sodium or potassium prepared by melting together 
litharge, PbO, potassium carbonate, K2COs3, and silica, SiOz. 
The name comes from the former use of crushed flints for the 
silica. Flint glass has a higher index of refraction but also a 
relatively greater dispersive power than crown glass, and the two 
varieties of glass are used together for achromatic lenses and for 
direct vision spectroscopes. It melts easily. Strass or paste 
is a heavy lead glass with a high index of refraction, used in 
making imitations of diamonds and other precious stones. 
Bohemian glass is a silicate of potassium and calcium having a 
high melting point. It is also less soluble in water than ordinary 
glass and is used for combustion tubing and for beakers and 
flasks used in the laboratory. It is often called hard glass. It 
has been largely replaced by various borosilicate glasses. 
“Jena” glass, “ Resistanz”’ glass and “ Non-sol” glass are 
borosilicates containing a little zinc. They are much less sol- 
uble in water than the ordinary glasses, and some of them soften 
at much higher temperatures. These properties render them 
suitable for special uses in chemical laboratories, especially for 
combustion tubing, for beakers and flasks for use in quantitative 
analysis and for test tubes for bacteriological cultures. Durax 
glass is a variety especially resistant to alkaline solutions. For 
the manufacture of thermometers several borosilicate glasses 
are made which give a much smaller depression of the zero point 
after use at high temperatures than is the case with thermometers 
made from ordinary glass. Some of these glasses also give 
thermometers which correspond closely with the hydrogen 
thermometer at high temperatures. 

Strontium, Sr, 87.63. Occurrence. As the most important 
natural compounds of calcium are the carbonate and sulfate, so 


468 A TEXTBOOK OF CHEMISTRY 


strontium and barium are found chiefly as carbonates and 
‘sulfates. The sulfates are much less soluble than calcium sul- 
fate. Strontium carbonate, SrCOs, is called strontianite, and 
strontium sulfate, SrSO,, celestite. The latter name is given 
because the mineral is often of a light blue color, but the pure 
sulfate is white. Water dissolves about 20 times as much cal- 
cium sulfate as it does of strontium sulfate. 

* Strontium Hydroxide, Sr(OH)>2.8 H,O, is more soluble than 
calcium hydroxide and is much more easily soluble in hot than in 
cold water. It forms a difficultly soluble compound with cane 
sugar, Cy2H».Oi1, and is sometimes used to recover sugar from 
the molasses of the beet sugar manufacture. 

Strontium Nitrate, Sr(NQO3)s, is used, mixed with sulfur, 
charcoal and potassium chlorate or nitrate for red lights in fire- 
works. 

Barium, Ba, 137.37. Occurrence. Barium is found in nature 
as the carbonate, BaCOs, called witherite, and the sulfate, BaSO,, © 
called barite. The latter is more common. 

Barium Oxide, BaO. Barium carbonate is very much more 
stable than calcium carbonate. The dissociation pressure of 
calcium carbonate reaches atmospheric pressure at 898°, but that 
of barium carbonate is equal to one atmosphere at 1350° (Finket- 
stein, Ber. 39, 1588 (1906)). Even at this temperature the 
decomposition, at first, appears to givea basic carbonate, prob- 
ably BaCO;.BaO, and a temperature of 1450° is required for 
the decomposition of this compound at atmospheric pressure. 
It is impracticable, therefore, to prepare barium oxide by the 
direct decomposition of the carbonate. If the carbonate is 
mixed with charcoal, however, the carbon dioxide formed by 
the dissociation of the barium carbonate will at once react with 
the carbon to form carbon monoxide. As a result of the two 


equilibria : BaCO; = BaO + COs 
CO, + C3Z2CO 


the constant removal of the carbon dioxide by means of the 
second reaction makes it possible to prepare barium oxide at a 


ALKALI-EARTH METALS: BARIUM 469 


temperature at which the dissociation pressure is small and the 
preparation is carried out, technically, by this method. 

Barium oxide may also be prepared by the decomposition of 
the nitrate, Ba(NO3)2, which occurs at a much lower temperature 
than that for the decomposition of the carbonate. 

Barium oxide combines directly with water to form the hydrox- 
ide, Ba(OH)s, the reaction being accompanied by considerable 
evolution of heat. It is used, chiefly, for the preparation of the 
peroxide, BaOz. It is also used in the laboratory as a powerful 
dehydrating agent, as in the preparation of absolute alcohol. 

Barium Peroxide, BaO:. ‘The dissociation pressure for the 
reaction, 2 BaO + On = 2 BaOx, is: 


Temperature 525. 670° 1305 hide 790° 
Pressureinmm. 20 80 260 510 670 


When we remember that the partial pressure of oxygen in the 
air is 21 per cent of 760 mm., or 160 mm., it is evident that if 
air is passed over barium oxide at a temperature of 670°, oxygen 
will be absorbed with the formation of the peroxide. On the 
other hand, if the peroxide is heated to about 800° under atmos- 
pheric pressure, the peroxide will be decomposed with evolution 
of oxygen. Again, if barium oxide at 800° is subjected to the 
action of air under a pressure of about five atmospheres, so that 
the partial pressure of the oxygen is a little more than one 
atmosphere, oxygen will be absorbed, and on lowering the pres- 
sure to atmospheric pressure oxygen will be evolved without 
any change in temperature. Both methods have been used for 
the technical preparation of oxygen. A little water vapor must 
be present to catalyze the reaction, and the carbon dioxide of 
the air must be carefully removed. 

Barium peroxide is also prepared and used for the manu- 
facture of hydrogen peroxide, H2,O2. It forms a_ hydrate, 
BaOz.8 H2O, which is difficultly soluble and which is formed on 
treating the anhydrous barium peroxide with water. It is also 
formed by precipitating an ice-cold solution of barium chloride, 
BaCl., with a cold solution of sodium peroxide, Na2Oz. 


470 A TEXTBOOK OF CHEMISTRY 


Barium Hydroxide, Ba(OH)s.8 H:O, is much more easily 
soluble in water than calcium or strontium hydroxides. It is 
also much more easily soluble in hot than in cold water and can 
be easily recrystallized from hot water. It gives a precipitate of 
barium carbonate, BaCOs, with carbon dioxide, and is a very 
sensitive reagent for the qualitative detection or quantitative 
determination of that gas. It is used to prepare an alkali 
solution which is free from carbonate, but such a solution must, 
of course, be carefully protected from the carbon dioxide of the 
alr. 

Barium Chloride, BaCle.2 H.O, is the most -common sol- 
uble salt of barium. It is used especially for the detection and 
quantitative determination of sulfates. 

* Barium Nitrate, Ba(NOs3)e, is sometimes used for the detec- 
tion of sulfates in solution containing silver or other metals which 
form insoluble chlorides. | 

* Barium Sulfide, BaS, is formed by heating barium sulfate, 
the most plentiful natural source of barium compounds, with 


charcoal : BaSO, + 4C = BaS + 4CO 


Barium sulfide dissolves in hydrochloric acid with the forma- 
tion of the chloride or in nitric acid with the formation of the 
nitrate, methods which were formerly used for the preparation 
of these salts from the insoluble sulfate. 

Barium Sulfate, BaSO, The mineral barite is sufficiently 
abundant to form a cheap source for barium compounds. It is 
also sometimes used in a finely pulverized form as an adulterant 
for white lead. It is crystalline and the particles are much more 
transparent than those of white lead, so that paints containing 
it have less covering power than those made from pure lead 
compounds. It is used in many mixed paints, especially in 
“lithopone,”’ which is a mixture of barium sulfate, BaSOu, and 
zinc sulfide, ZnS, obtained by precipitating a solution of zinc 
sulfate with barium sulfide, BaS. Lithopone is not blackened 
by hydrogen sulfide and is more suitable than white lead for 
places where that gas is liable to be present. 


RADIUM 471 


Barium sulfate requires about 400,000 parts of water for its 
solution and it is only slightly more soluble in dilute acids. It 
is very much used for this reason for the detection and estima- 
tion of sulfates. It hasa very marked tendency to form solid 
solutions or mixed crystals with other sulfates, barium chloride 
or other compounds which may be present. The presence of 
these foreign substances may lead to rather serious errors in the 
determination of sulfates, unless great care is used in following 
proper methods of manipulation. 

Barium sulfate is appreciably soluble in solutions containing 
considerable hydrochloric acid and dissolves rather easily in 
concentrated sulfuric acid. 

Flame Colors for Calcium, Strontium, and Barium. Calcium 
compounds impart a brick red color to the Bunsen flame, stron- 
tium compounds a bright red and barium compounds a green 
color. The spectra show bright lines and bands which are easily 
distinguished, even with the simplest forms of spectroscopes. 

Radium (Ra, 225.95). About 1878 Sir William Crookes dis- 
covered that in the discharge of electricity through a highly 
rarefied gas rays are shot out at right angles to the surface of the 
cathode and produce a beautiful green fluorescence at the point 
where they strike the containing tube. About twenty years 
later it was shown that these rays consist of electrons traveling 
with velocities approaching that of light. In 1895 Rontgen, 
partly by accident, discovered that rays, afterwards called 
Rontgen rays, emanate from the glass at the point of fluores- 
cence. These rays penetrate paper, wood and some other ob- 
jects, which are opaque to ordinary light, and affect a photo- 
graphic plate placed behind such screens.. The rays are, however, 
intercepted by metals or substances containing compounds of the 
metals, as, for instance, by bones. The opacity of various sub- 
stances is closely proportional to their density. Very shortly 
after, in 1896, Becquerel, in Paris, discovered that minerals con- 
taining uranium have the property of emitting penetrating radi- 
ations resembling the Rontgen rays in their effect on a photo- 
graphic plate protected from ordinary light rays by a screen of 


472 A TEXTBOOK OF CHEMISTRY 


black paper. Further study by Madame and Monsieur Curie 
led to the discovery in uranium ores of a new element, radium. 
This belongs to the calcium-barium family, and forms a sulfate 
which is much less soluble than barium sulfate and which is made 
use of in separating radium from other elements. Radium and 
its compounds exhibit the following remarkable properties : 

1. It affects a photographic plate through black paper and 
will cause the fluorescence of zinc sulfide and some other com- 
pounds exposed to the action of its rays. 

2. It causes the ionization of air, that is, the separation of the 
molecules of the air into positive and negative, charged particles, 
causing the air to become a conductor of electricity. A gold-leaf 
electroscope, which will remain charged for a long time in ordi- 
nary air, is rapidly discharged and the leaves fall together in air 
which has been exposed to the action of radium. The rate of 
discharge furnishes a quite accurate measure of the quantity of 
radioactive substances present, and the measurement of this rate 
is the most important method used in the study of radioactive 
elements. 

3. Radium continually evolves heat. One gram of the ele- 
ment gives out 132 small calories per hour. ‘This phenomenon 
is independent of the temperature or of the form in which the 
element is combined or of any other conditions which can be 
controlled. 

Disintegration of Atoms. In 1902-1903 Professor Rutherford, 
then at McGill University in Montreal, published a series of 
papers in which he proposed the hypothesis that the atoms of 
radioactive elements disintegrate more or less rapidly, breaking 
down with the formation of other elements. In the disintegra- 
tion, portions of the atom are shot out from it with tremendous 
velocity. Some sort of potential or kinetic energy within the 
atom is liberated in this manner and manifests itself as heat 
energy, and this explains the heat evolved by radioactive ele- 
ments. Incidentally this makes it probable that all atoms are 
complex in their internal structure and are storehouses of 
immense quantities of energy. The particles shot out by the 


RADIUM: DISINTEGRATION THEORY 473 


disintegrating atoms seem to tear apart molecules which they 
strike, separating them into charged particles or ions, and in this 
manner air or other gases which are exposed to the action of 
radioactive substances become conductors of electricity as has 
been mentioned above. 

Rutherford’s disintegration theory was very strongly sup- 
ported when Soddy, who began work with Rutherford, demon- 
strated, while working with Sir William Ramsay in London, that 
helium is one of the disintegration products of radium. The 
theory is now accepted, at least as a working hypothesis, by all 
investigators in this field. A few chemists, on account of the 
disintegration of radium, have contended that it is not properly 
called an element but should be classed as a chemical compound. 
But radium finds its place in the periodic system and has all of 
the other properties which usually characterize an element. 
Moreover, there is good reason to believe that radium is, in turn, 
a disintegration product of uranium. If radium is a compound, 
uranium is a compound also. It seems better, therefore, to 
revise our definition of an element and accept the notion that 
the atoms of some elements, and possibly of all, may disintegrate 
with the formation of other atoms. The relation between an 
atom and its disintegration products is evidently very different 
from the relation between a compound and the elements of which 
it is composed. 

Nature of the Radiations from Radioactive Substances. 
Four kinds of rays have been distinguished as emanating from 
radioactive substances. The first kind, called e-rays, have been 
identified as atoms of helium, carrying a double positive charge. 
As they are charged particles moving with a high velocity, they 
are slightly deflected by a strong magnetic field. The f-rays 
appear to be identical with the cathode rays of the Crookes tube, 
i.e. they are electrons moving with varying velocities, the 
velocity sometimes approaching that of light. As the charge 
is negative, they are deflected by a magnetic field in a direction 
opposite to that of the a-rays; and as the mass is very much 
smaller in proportion to the charge, the deflection is much 


A74 A TEXTBOOK OF CHEMISTRY 


greater. The y-rays seem to be like the X-rays or Rontgen 
rays and are probably of the nature of ether waves. They will 
penetrate a very much greater thickness of metal than the 
B-rays and these in turn are much more penetrating than the 
a-rays. Some of the radioactive elements give out all three 
kinds of rays, others give only one kind and still others two kinds. 
The 6-rays are electrons moving much more slowly than those 
which form the B-rays. They may be given out from the sur- 
face of radioactive material, or may be emitted from any sub- 
stance traversed by «&-rays, whether solid or gaseous. The 
nature of the rays furnishes one of the most important means of 
identifying different elements. 

The Life of an Element. In accordance with the disintegra- 
tion theory some of the atoms of radium are constantly decom- 
posing into helium and another substance which was at first 
called radium emanation, but which has been characterized by 
Sir William Ramsay as an element of the argon family, the gas 
radon. This has an atomic weight of 222.4, the difference in 
weight between an atom of radium and an atom of radon being 
the weight of an atom of helium. The'rate of the decomposition 
of radium has been determined by measuring the amount of 
helium given in a number of weeks or months by a given weight 
of radium. The result of the measurement was that one half 
of a given quantity of radium would disintegrate in 1700 years. 
The rate of disintegration of radon, on the other hand, is so rapid 
that one half of a given quantity of the element will disintegrate 
in 3.8 days. The rate of decay is measured in this case by 
measuring the rate at which the radioactive effect on the elec- 
trometer decreases. The disintegration of uranium to form the 
first of a series of 5elements which are supposed to stand between 
uranium (at. wt. 238.2) and radium (at. wt. 226) takes place 
so slowly that one half of a given quantity of the element would 
decompose in 5,000,000,000 years. These periods, which are 
called the “half-life periods” of the elements, form one of the 
most important characteristics of the radioactive elements. It 
is evident from the periods which have been given for uranium, 


ISOTOPES 475 


radium, and radon that the amount of radium which can exist 
at a given time must be very small in comparison with the quan- 
tity of uranium in the world, and that the amount of niton must 
always be very small in proportion to the amount of the radium 
from which it is generated. For this reason Sir William Ramsay 
was compelled to establish the density and other properties of 
radon by working with a few cubic millimeters of the gas. He 
weighed the gas with a microbalance, designed for the purpose, 
with which it was possible to weigh spo/o00 of a milligram. 

Shift in the Periodic System of the Disintegration Product of 
an Element. The atomic number of an element has been defined 
as either the excess of the protons over the electrons in the 
nucleus or the number of electrons surrounding the nucleus. 
When a radium atom disintegrates it shoots out an alpha par- 
ticle carrying two positive charges, as stated above. This is 
evidently the nucleus of a helium atom, which consists of four 
protons and two electrons. This must leave the nucleus of an 
atom of radon, which contains two positive charges less than 
the nucleus of the radium atom. The nucleus of radon can, 
therefore, hold two fewer electrons about it, than the nucleus 
of radium. This must allow the two valence electrons of the 
radium to escape. The loss of the alpha particle and of the 
two electrons will leave an atom of radon. Ina similar manner 
the loss of an alpha particle in the disintegration of any radio- 
active element must cause a shift to the left of two places in 
the Periodic System. 

The loss of an electron, or 8-particle, from the outer shell of 
an atom must leave the nucleus with an extra positive charge. 
This can be balanced by sharing one more valence electron with 
some other atom. This corresponds to a shift of one place to 
the right in the Periodic System. Shifts of this sort have been 
observed for all the radioactive elements and furnish a very 
satisfactory confirmation of the theory of the structure of atoms 
(p. 188). 

Other Radioactive Elements. Isotopes. Nearly 40 radio- 
active elements have been identified by the nature of the rays 


476 A TEXTBOOK OF CHEMISTRY 


which they emit and their life as shown by the decrease of the 
intensity of their radiations with the time. Only uranium, 
thorium, radium and radon have, thus far, been isolated and 
their atomic weights determined by the usual methods. The 
others are known only by their radioactive properties. 

There are three well-defined series, with uranium, thorium 
and actinium, having atomic numbers 92, 90 and 89, as the 
first element of the respective series. The last disintegration 
product of each series seems to be lead with an atomic number 
of 82. The lead formed by the disintegration of the uranium 
series has an atomic weight approaching 206, while the atomic 
weight of ordinary lead is 207.2. The two kinds of lead are 
exactly alike in their chemical properties and very closely alike 
in their spectra. It is supposed that they differ in the number 
of protons and electrons in their nuclei but that they are identi- 
cal in the number of protons in excess of electrons and in the 
number of electrons in their shells. Atoms related in this way 
are called zsotopes. 

The 40 radioactive elements mentioned above all have atomic 
numbers between 82 and 92. It is evident that there must be 
about four isotopes for each atomic number. But while radio- 
active isotopes are identical in their chemical properties their 
half-life may differ from a minute fraction of a second to several 
years. It would seem that radioactive isotopes may differ 
greatly in the structure of their nuclei or in the arrangements 
of the electrons about the nuclei. The half-life periods of the 
majority of the radioactive elements are very short. 

Chemical Action of the Rays. If radium or niton is left in 
contact with water, the rays which they emit cause dissociation 
of some of the water into oxygen and hydrogen. Some hydro- 
gen peroxide is also formed. A glass tube containing a com- 
pound of radium soon assumes a violet or brown color. The 
tremendous velocity with which a-particles are expelled from 
radium or radon gives a unique and powerful form of energy, 
and atoms of other elements have been broken into pieces by 
this means (Rutherford). 


RADIOCHEMISTRY 477 


Radiochemistry in Relation to Geology and Medicine. Radio- 
active elements are very widely diffused in the rocks of the 
earth; and while the proportion of such elements is very small, 
it has been shown that the total amount present in the crust of 
the earth is sufficient to account for the increasing temperature 
which is observed in deep wells and mines and in tunnels. 
Indeed it would seem that the proportion of radioactive ele- 
ments must be smaller at very great depths than it is near the 
surface. This discovery has thrown very grave doubts on esti- 
mates formerly made of the life of the earth, which were based 
on the supposition that the earth has cooled down from a molten 
condition. 

The rays emitted from radium and other radioactive elements 
are fatal to bacteria. They also may produce severe burns 
somewhat resembling sunburn. They have been used with 
some success in the treatment of cancer, ulcers, lupus, ete. 
Some mineral waters in which chemical analysis has formerly 
shown no peculiar curative substances have been found to be 
radioactive, and it seems possible that beneficial results may be 
obtained by the use of such waters. 


EXERCISES 


1. How many grams of water will be required, theoretically, to 
convert a pound (453 g.) of plaster of Paris into gypsum ? 


2. How much lime can be obtained from a kilogram of marble ? 
3. The heat of formation: 
CaO + CO, > CaCO; 


is 42,520 small calories per gram molecule. How much coal having the 
same heat of combustion as carbon would be required, theoretically, 
to prepare one kilogram of lime? See p. 27 for the heat of combustion 
of carbon. 


4. The heat of formation: 
BaO + CO, — BaCOs; 


is 62,220 small calories per gram molecule. How much coal having the 
same heat of combustion as carbon will be required to give one kilogram 
of barium oxide on the basis of the reaction: 


BaCO; + C = BaO + 2 CO? 


CHAPTER XXVII 


ALTERNATE METALS OF GROUP II. MAGNESIUM, ZINC, 
CADMIUM AND MERCURY 


In its occurrence and in the properties of its most common 
compounds, magnesium resembles calcium rather than zinc, but 
in the metallic form the resemblance to zine is more marked. 
Magnesium decomposes water slowly at the boiling point, while 
zinc and cadmium decompose it readily at a higher temperature. 
Mercury in many of its properties seems to be more closely 
related to copper and silver than to cadmium, and some authors 
formerly placed it in the first group. It does not decompose 
water at any temperature and it forms compounds in which it is 
apparently univalent as well as those in which it is bivalent. 
Magnesium, zine and cadmium are always bivalent. 

Magnesium, Mg, 24.32, is found as the carbonate, magnesite, 
MegCOs, as a double carbonate of calcium and magnesium, 
dolomite, CaCO3.MgCOs, as the double chloride with potassium, 
carnallite, KCl.MgCle.6 H2O, and as a principal constituent of 
many silicates, especially tale or soapstone, Mgs3H2(SiOs3),, 
serpentine, Mg3Si207.2 H2O, and meerschaum, MgoSi303.2 H.O. 
The sulfate, Epsom salts, MgSO..7 H,0, and the chloride, 
MegCly.6 H2O, are also found in many natural waters. 

Preparation, Properties. Metallic magnesium is obtained 
by the electrolysis of fused carnallite, MgCl..KCl, from which 
the water of hydration has been expelled by heat. It is a 
silver-white, very light metal, having a specific gravity of only 
1.75, slightly lower than that of beryllium. It melts at 651° 
and boils at about 1100°. Magnesium wire or ribbon burns in 
the air with a very intense white light that is particularly rich 


in the more refrangible rays, which affect the photographic plate. 
478 


GROUP II: MAGNESIUM 479 


For this reason powdered magnesium is the effective constituent 
of flash-hght powders. It is estimated that 10 per cent of the 
energy of burning magnesium appears as light, a very much 
larger per cent than is secured by any ordinary illuminant. 
The temperature of burning magnesium is not, however, very 
high — only about 1340°. Magnesium tarnishes only slightly, 
if at all, in dry air and very slowly in moist air, so that it can 
be kept indefinitely without special precautions. 

Magnesium is used in the laboratory as a powerful reducing 
agent for the preparation of silicon and boron. It is also used 
in a great variety of syntheses of organic compounds. 

Magnesium Oxide, MgO, is most easily prepared by heating 
magnesium carbonate, MgCOs:, which decomposes at a much 
lower temperature than calcium carbonate. It is a light, white 
powder often called magnesia usta, or burnt magnesia. It is 
infusible in any ordinary furnace, but may be volatilized in the 
electric furnace. It is used for crucibles and for some forms of 
apparatus which must withstand extremely high temperatures. 
It is used as a basic lining for metallurgical furnaces, especially 
for the basic process for steel (p. 548). 

Magnesium Hydroxide, Mg(OH)s, is obtained as a white 
precipitate on the addition of sodium hydroxide, NaOH, or barium 
hydroxide, Ba(OH)s, to a solution containing almost any soluble 
salt of magnesium. It is much less soluble than calcium hydrox- 
ide, dissolving in about 6500 parts of water. In spite of this 
difficult solubility, however, it is not precipitated in the presence 
of ammonium salts. This is due probably to two reasons: 
(1) in the presence of an ammonium salt the ionization, 


NH.OH = NH,* + OH- 


is repressed by the ammonium ions, NH4', of the salt, and the 
concentration of the hydroxide ions, OH~, is low in such a solu- 
tion; and (2) because magnesium hydroxide as a unibivalent com- 
pound very probably forms intermediate ions, MgOH* + OH, 
which interfere with the application of the ordinary law of the 
solubility product (p. 377). 


480 A TEXTBOOK OF CHEMISTRY 


Magnesium hydroxide is easily decomposed by heat into 
magnesium oxide and water. 

Magnesium Chloride, MgCl».6 H2O, crystallizes from a con-_ 
centrated solution of the salt. It is very easily soluble in 
water. When an attempt is made to drive out the water of the 
salt by heating it, both water and hydrochloric acid escape, and 
a mixture of variable composition, containing chiefly mag- 
nesium oxide, finally remains : 


MgCl, + H.O = MgO + 2 HCl 


The process has been used to a limited extent as a basis for the 
preparation of hydrochloric acid and chlorine. 

Magnesium Ammonium Chloride, MgNH,Cl3.6 H.O. This 
salt is easily prepared by crystallizing from water a mixture of 
equimolecular amounts of magnesium chloride and ammonium 
chloride. The water of hydration may be expelled with very 
little loss of hydrochloric acid, and on heating the anhydrous 
salt to a slightly higher temperature the ammonium chloride 
dissociates and escapes, leaving anhydrous magnesium chloride 
behind. | 

Magnesium Sulfate, MgSO,.7 H.O, or Epsom Salts, is found 
in some mineral waters used for their medicinal properties, es- 
pecially in Hunyadi water, in which it is associated with sodium 
sulfate, NaeSOu. 

* Magnesium Sulfide, MgS, may be prepared by heating a 
mixture of magnesium and sulfur. It is decomposed by water, 
giving magnesium hydroxide and hydrogen sulfide: 


MgS + 2 HOH = Mg(OH). + HS 


Both the insolubility of the hydroxide and the volatility of the 
hydrogen sulfide contribute to cause the reaction to go to com- 
pletion. 

* Magnesium Ammonium Phosphate, MgNH.PO.,.6 H.O, is 
a difficultly soluble salt which is formed when solutions contain- 
ing magnesium, ammonium and a soluble phosphate are brought 
together. It is used in analytical chemistry for the determina- 
tion of both magnesium and phosphoric acid. A precipitate 


GROUP II: ZINC 481 


having the exact composition represented by the formula can 
be obtained only by securing exactly the right conditions as 
regards the concentration of the various solutions employed. 
When heated, the compound decomposes quantitatively with the 
formation of magnesium pyrophosphate, Mg2P.2Ov. 

Zinc, Zn, 65.38. Occurrence. Zinc is found in nature as the 
sulfide, sphalerite, ZnS, the carbonate, smithsonite, ZnCOs3, the 
silicates, willemite, Zn2SiO4u, and calamine, H2ZnSiOu, and in the 
mineral franklinite, which consists of oxides of iron, zine and 
manganese. Franklinite is an important ore in New Jersey and 
is used as a source of zinc, manganese and iron. 

Metallurgy. The sulfide of zinc is converted into the oxide 
by roasting it, that is by heating it in a furnace with free access 
of air: 


ZnS + 30 = ZnO + SO, 


The oxide, obtained in this way or by heating the carbonate, 
is mixed with coal and heated to a high temperature in an earth- 
enware retort having a receiver luted to it with clay. The zine 
oxide is reduced, and the zinc, which boils at 925°, distils over. 
Zinc melts at 419.4°. 

Impure zinc dissolves easily in hydrochloric or in dilute sul- 
furic acid, with evolution of hydrogen. Pure zine, which can 
be obtained by distillation in a vacuum, is attacked very slowly 
or not at all by these acids, but dissolves readily in contact with 
platinum. It has been pointed out that these facts indicate that 
the solution is always associated with electrical phenomena. 
Impure zine which is covered by a thin film of amalgam, giving 
it a homogeneous surface, is also not attacked by the dilute acids, 
and such amalgamated zinc is used in electric batteries. Zine 
has a specific gravity of 6.9. | 

Uses. Galvanized Iron. Metallic zinc is used chiefly in brass, 
an alloy of the metal with about twice its weight of copper, and 
as a coating for iron to protect it from rusting. It is also a con- 
stituent of many of the bronzes, and is used as the metal which is 
dissolved or corroded in most forms of primary electrical batter- 


482 A TEXTBOOK OF CHEMISTRY 


ies, especially in the gravity cell used in telegraphy and in the 
so-called “‘ dry ”’ batteries. 

“Galvanized ” iron is prepared by dipping carefully cleaned 
sheet iron or other iron or steel articles in melted zinc. The 
value of the coating depends on two properties: first, zinc is 
electropositive with reference to iron and when the two metals 
are in contact with each other and also in contact with an elec- 
trolyte the zinc is attacked and the iron is protected ; second, the 
action upon the zine causes the formation of a very thin, 
coherent coating of zine oxide or hydroxide, which is practically 
insoluble in water and protects the zine from further action. 
A small amount of zine passes into solution, however, and this 
may be increased very considerably in the presence of even 
weak acids. As zinc salts are poisonous, pails or dishes of gal- 
vanized iron are not suitable for culinary use. 

The coating of zinc has been sometimes applied to the iron by 
an electrolytic method, and the term “ galvanized iron” came 
from this method of manufacture. 

Sherardized Iron. A new process for coating iron with zine 
has been developed by Sherard Cowper-Cowles. The articles to 
be coated are heated with zinc dust in iron drums at 500°-600° 
for thirty minutes to several hours, according to the thickness of 
the coating desired. The process is somewhat analogous to the 
manufacture of cementation steel. See Johnson and Woolrich, 
Trans. Am. Electrochem. Soc. 21, 561 (1912). Cowper-Cowles, 
Electrochem. and Met. Ind. 6, 189 (1908). 

Zinc Oxide, ZnO, is prepared by burning the vapors of metallic 
zinc. It is a white powder and gives with linseed oil an excellent 
pigment, which has the advantage of not being blackened by 
hydrogen sulfide because zine sulfide is also white. Zine oxide 
is often used with phosphoric acid for a cement in dental work. 
The two substances combine to form a basic zine phosphate 
which sets to a hard mass that adheres strongly to the surfaces 
with which it is in contact. Zinc oxide is yellow when hot, but 
turns white again on cooling, a property used for the detection 
of zinc compounds in blowpipe reactions. 


GROUP II: ZINC, CADMIUM 483 


Zinc Hydroxide, Zn(OH)2, forms as a white precipitate on 
adding a soluble hydroxide to a solution of a zine salt. It dis- 
solves in an excess of sodium hydroxide or potassium hydroxide, 
forming sodium zincate, Na2ZnOx, or potassium zincate, KeZnOxz. 
In forming these compounds zinc hydroxide seems to act as an 
acid, which might be called zincic acid, and the formula might be 
written HsZnO.. Toward acids, however, zinc hydroxide con- 
ducts itself as a true hydroxide or base. Compounds which 
exhibit a dual nature of this sort, acting in some conditions as 
acids and in others as bases, are said to be amphoteric. 

Zinc. Chloride, ZnCl,. An aqueous solution of this salt is 
easily prepared by dissolving metallic zine or zinc oxide in hydro- 
chloric acid. Unlike magnesium chloride, the solution loses 
only a small amount of hydrochloric acid when heated to a high 
temperature to expel the water. The pure, anhydrous salt 
melts at 290°-297° and boils at 730°. When boiled in an iron 
tube, it furnishes an easy means of securing a constant, rather 
high temperature and has been found useful for this purpose. 
Zine chloride is used in the treatment of wooden ties to prevent 
decay. 

Zinc Sulfate, ZnSO,.7 H.O, or White Vitriol. Anhydrous 
zine sulfate, ZnSO, can be prepared by roasting the sulfide, ZnS, 
at a moderate temperature. At a higher temperature the sul- 
fide roasts to the oxide and sulfur dioxide. The hydrate, 
ZnSO,.7 HO, forms rhombic crystals and is easily soluble. 

Zinc Sulfide, ZnS, forms as a white precipitate when hydrogen 
sulfide is passed into a neutral or alkaline solution of a zinc salt. 
The precipitate is formed even in slightly acid solutions, and care- 
ful attention must be paid to the amount and character of the 
acid present if a separation from other metals is desired. In 
the presence of sulfuric acid which is weaker than fifth normal 
(about 1 per cent) the sulfide will be precipitated. 

Cadmium (Cd, 112.41). Many zinc ores contain a small 
amount of cadmium. As the boiling point of cadmium (785°) 
is considerably lower than that of zinc (925°), the former distills 
over first in the preparation of zine, and by collecting these por- 


484 A TEXTBOOK OF CHEMISTRY 


tions and subjecting them to fractional distillation, nearly pure 
cadmium can be prepared. 

Metallic cadmium closely resembles zine in appearance and 
in many of its properties. It melts at 320.9°, boils at 785° and has 
a specific gravity of 8.65. It was formerly used in amalgams 
for filling teeth, but other amalgams are now considered more 
suitable. It is a constituent of Wood’s metal and of the easily 
fusible alloys used for safety fuses in electrical circuits and for 
automatic sprinklers used for protection against fire. 

* Cadmium Hydroxide, Cd(OH)p, is easily obtained as a white 
precipitate. It dissolves easily in acids but does not dissolve in 
solutions of sodium or potassium hydroxides, as zinc hydroxide 
does. It is decomposed when heated, giving cadmium oxide, 
CdO, as a brown powder. 

* Cadmium Sulfate, 3 CdSO,.8 H.O, is an easily soluble salt 
used in the Weston standard cells, which are the most satisfac- 
tory primary standard for the measurement of electromotive 
force. 

Cadmium Sulfide, CdS, forms as a yellow precipitate in solu- 
tions of cadmium salts which do not contain too much free acid 
or too much of salts which interfere with the precipitation. 
From a solution containing both zine and cadmium in which 
sulfuric acid is present and the concentration of the hydrogen ion 
is between fifth normal and twice normal and other interfering 
salts or acids are absent, the precipitation of the cadmium is 
practically complete, while only a small amount of the zine will 
come down. For a complete separation, however, the cadmium 
sulfide must be dissolved and reprecipitated. In hydrochloric 
acid stronger than 0.3 normal cadmium sulfide is not completely 
precipitated. Cadmium sulfide dissolves readily in boiling, 
dilute sulfuric acid, but is insoluble in a solution of potassium 
cyanide, KCN. 

Mercury, Hg, 200.6. Occurrence. Metallurgy. From the 
positions of the elements of Group II in the electromotive 
series (p. 436) mercury is the only element of the group which 
could appear in the free state in nature. It is occasionally found 


GROUP IL: MERCURY 485 


in small globules disseminated in porous rocks. Mercury occurs 
chiefly, however, in the form of the native sulfide, cinnabar, HgS, 
a brilliant red mineral, when pure. When ores containing cinna- 
bar are roasted by heating in a current of air, the sulfur burns to 
sulfur dioxide while the mercury distills and is condensed in 
long flues where the vapors must be very thoroughly cooled to 


prevent loss: HeS + O. = He + SOz 


Mercury may also be obtained by mixing the sulfide with lime 
or with iron and distilling : 


2 CaO + 2 HeS = 2 CaS + 2Hg+ QO, 
Fe + HgS = FeS + Hg 


Mercury can be purified by allowing it to fall in very minute 
globules, through a chamois skin tied over the end of a glass 
funnel, into dilute nitric acid contained in a tube 2 meters long. 
The tube is drawn out and bent upward at the bottom so that a 
short column of mercury in the overflow tube balances the col- 
umn of nitric acid. The nitric acid dissolves zinc, arsenic, lead 
and nearly all of the other metals likely to be present. 

Mercury may also be separated from nearly all other metals 
by distillation under diminished pressure. If a very little air 
is allowed to pass through the mercury by means of a very fine, 
hairlike, capillary tube, troublesome bumping of the mercury 
can be avoided and zinc and some other metals are oxidized, 
giving purer mercury than if the distillation is carried out in 
the absence of air (Hulett). 

Properties and Uses. Mercury isa heavy, mobile liquid, with 


a density, at i = 13.5956 or at a err sode3.. ctreeresiat 
— 38.70°. Its freezing point on a Fahrenheit thermometer is 
— 37.7°. It will be noticed that the scales of the Centigrade 
and Fahrenheit thermometers approach very closely together 
at this temperature. Mercury boils at 357°. Its critical tem- 
perature is about 1275°, and its critical pressure is calculated as 


about 675 atmospheres, an extraordinarily high value. (Menzies, 


486 A TEXTBOOK OF CHEMISTRY 


J. Am. Chem. Soc., Sept., 1918. Konigsberger, Chem. Ztg. 138, 
1321 (1913).) Mercury oxidizes slowly to red mercuric oxide, HgO, 
when heated to its boiling point in the air. (See Lavoisier’s 
experiment, p. 19.) Mercury dissolves very slowly to mercur- 
ous nitrate, HgNOs, in dilute nitric acid, with evolution of nitric 
oxide, NO. It is insoluble in hydrochloric acid, but is converted 
into mercurous sulfate, HgeSO., by hot, concentrated sulfuric 
acid, with evolution of sulfur dioxide. 

Mercury is used in the amalgamation processes for the re- 
covery of gold and silver (p. 441), in making thermometers, for 
barometers and manometers, in mercury air pumps and in the 
collection and measurement of gases. Its advantages over all 
other substances used in thermometers are, especially, that it 
does not wet or attack the glass, that it is liquid over a wide 
range of temperature, including the common range of air tem- 
peratures, and that its rate of expansion is very uniform. Its 
coefficient of expansion between 0° and 100° is so nearly constant 
that a mercury thermometer graduated in equal degrees does 
not differ from the standard hydrogen scale by more than 0.2° 
at any point between these temperatures. Mercury thermom- 
eters cannot, of course, be used at temperatures below — 39° 
and ordinary thermometers cannot be used above 300° — 
indeed, the thread of an ordinary thermometer will usually break 
before that temperature is reached. By filling the space 
above the mercury with nitrogen under pressure, however, 
thermometers graduated to 460° are made, and by filling the 
space with carbon dioxide the range has been carried to 550° or 
above. There is likely to be a large zero-point correction for 
such thermometers, and they must be carefully treated, if 
accurate results are required. ‘The stem correction is also 
large, unless the whole thermometer is immersed in the sub- 
stance whose temperature is to be measured. 

Amalgams. ‘The alloys of mercury are called amalgams. 
Many metals, such as sodium, potassium, copper, silver, gold, 
zinc, cadmium, tin and lead, dissolve in or alloy with mercury 
in all proportions or give amalgams having a wide range in their 


GROUP II: AMALGAMS 487 


composition. Other metals, as iron and platinum, dissolve in 
mercury to only a trifling extent or not at all. The amalgams of 
gold and silver are used to separate these metals from large 
masses of other substances mixed with them in their ores. 
Sodium amalgam is often used as a reducing agent, especially 
for organic compounds. Zinc amalgam and other amalgams 
may be used in the same way. An amalgam of tin was formerly 
used for the backs of mirrors, but has been replaced by a thin 
film of metallic silver in modern mirrors. An amalgam with 
silver and other metals is used for filling teeth. 

In many cases mercury combines with metals to form definite 
compounds. Such compounds are most easily identified by a 


v= ese ee eps 
AHHH 
200 | 


TEMPERATURES. 





-50° 
Bet rinOanl ogi) 20th SOU PELOS EB 60.160 i570) h:80,4 90.200 


FerCetlti}i99 99 0 70 60 50 40 30 2 10 0 


study of the freezing point curveof amalgams of varying composi- 
tion. In Fig. 102 the ordinates give the melting points and the 
abscissas give the composition of a series of amalgams of mercury 
with sodium. It will be seen from the figure that the addition 


A488 A TEXTBOOK OF CHEMISTRY 


of mercury to sodium lowers its melting point until a minimum 
is reached for an amalgam containing about 40 per cent of mer- 
cury and 60 per cent of sodium. This minimum is called a 
eutectic point. An amalgam of this composition melts at 21°. 
Further addition of mercury raises the melting point till a maxi- 
mum is reached for anamalgam containing 5.4 percent of sodium 
and 94.6 per cent of mercury, which melts at 346°. Further 
addition of mercury lowers the melting point till this would, 
undoubtedly, fall below the melting point of pure mercury. A 
compound having the formula NaHge would contain 5.43 per 
cent of sodium and 94.57 per cent of mercury. Evidently the 
amalgam of the highest melting point is a compound of this 
formula. The addition of either sodium or mercury to this 
compound lowers its melting point just as the addition of salt or 
any soluble substance lowers the melting point of ice. Changes 
in the direction of the curve at other points indicate that other 
compounds of mercury and sodium are present in some of the 
amalgams, but these details are not shown in the figure. See 
Kurnakow, Z. anorg. Chem. 23, 443 (1900). 

Compounds of Mercury. Mercury forms mercurous com- 
pounds, such as HgeO, HgeCle, Hg2SOz, in which it appears uni- 
valent, but in which it is probably really bivalent as expressed 

Hg—Cl 
by the graphical formula, | . It also forms mercuric com- 
Hg—Cl 
pounds, such as HgO, HgCle, HgSO,, in which it is clearly bi- 
valent. In the formation of these two classes of compounds and 
also in its conduct toward nitric, sulfuric and hydrochloric acids 
mercury resembles copper rather than zine or cadmium, and it 
has sometimes been classified under the first group of the Periodic 
System in place of gold. 

Mercurous Oxide, Hg.O, is formed as a black precipitate 
when a solution of sodium hydroxide is added to a solution of 
mercurous nitrate, HgNOs, or when calomel, HgeClo, is digested 
with a solution of sodium hydroxide. 

Mercuric Oxide, HgO, is formed slowly as a heavy red crys- 


GROUP II: MERCURY 489 


talline powder when mercury is heated to its boiling point in the 
air. Itis obtained more easily by heating the nitrate. A yellow 
precipitate having the same composition is formed on adding an 
alkali to a solution of a mercuric salt. 

Mercuric Sulfide, HgS. The mineral cinnabar, HgS, is a 
bright red compound. When hydrogen sulfide is passed into a 
solution of a mercuric salt a black mercuric sulfide of exactly 
the same composition is precipitated. By subliming the black 
sulfide, or by warming it with a solution of sodium sulfide, it 
can be converted into the red variety. The red form is used 
under the name of vermilion as a brilliant red pigment. When 
applied to iron or zinc, however, it is decomposed with libera- 
tion of metallic mercury. 

Mercurous Chloride, or Calomel, Hg2Cle, is prepared by sub- 
liming a mixture of mercuric chloride, HgClo, and mercury, or a 
mixture of mercuric sulfate, HgSO., salt and mercury. The 
crude product usually contains a little mercuric chloride, which is 
removed by treatment with alcohol, in which the mercurous 
chloride is insoluble while the mercuric chloride is easily soluble. 
The gram molecular volume of the vapor of mercurous chloride 
weighs about 236 grams, corresponding to the formula HgCl, 
but it has been shown that the vapor really consists of a mixture 
of mercuric chloride and mercury (HgCle + Hg) (Alex. Smith, 
J. Am. Chem. Soc. 32 1541 (1910)). 

From this it seems probable that the true formula of mercurous 
chloride is Hg2Cle. 

Calomel is used as a medicine. It is now usually administered 
in very small doses and mixed with sodium bicarbonate, NaHCOs, 
to render it less soluble in the acid gastric juice. In former times 
the careless administration of large doses sometimes caused sali- 
vation and other serious injuries to patients. 

Mercurous chloride is formed as a white precipitate on adding 
hydrochloric acid or a soluble chloride to a solution of mercurous 
nitrate or of some other soluble mercurous salt. 

Mercuric Chloride or Corrosive Sublimate, HgClo, is prepared 
by subliming a mixture of mercuric sulfate, HgSOu., and salt, 


490 A TEXTBOOK OF CHEMISTRY 


NaCl. It is a white, crystalline salt, which melts at 265° and 
boils at 307°. It is soluble in about 14 parts of cold water and 
more easily soluble in alcohol. It also dissolves in ether. When 
taken internally it is very poisonous. The best antidote is the 
white of an egg, with which it forms an insoluble compound. 
The solution in alcohol is sometimes used as a poison for insects. 
A dilute solution (usually 1: 1000) is much used as an antiseptic 
in surgery. From such a solution the mercury does not seem to 
be absorbed from a wound or through the skin. 

Mercuric Iodide, Hgls, is precipitated as a scarlet powder 
on adding a solution of potassium iodide, KI, to a solution of 
mercuric chloride. The precipitate dissolves in an excess of the 
potassium iodide, forming the complex salt, K.HgIy. Sodium 
hydroxide produces no precipitate in such a solution, evidently 
because it contains only a very small number of mercuric ions, 
Hg**. An alkaline solution prepared in this manner is used 
under the name of Nessler’s solution as an extremely sensitive 
reagent for ammonia. 

Mercurous Nitrate, HgNO3 or Hg2(NOs3)s, is formed by the 
solution of mercury in cold, dilute nitric aeid. It is hydrolyzed 
by water, giving a basic nitrate, Hge(OH)NOs, hence to secure 
a clear solution a little nitric acid must be added to carry the 
reversible reaction : 


to the left. To counteract the oxidation to mercuric nitrate, 
Hg(NOs3)e, by the oxygen of the air, some metallic mercury must 
be kept in contact with the solution : 


Hg(NOs3)o + Hg = Hg2(NQO3)o 


* Mercuric Nitrate, Hg(NO3)2.8 HO, is obtained by dissolv- 
ing mercury in warm, concentrated nitric acid. 

* Mercuric Cyanide, Hg(CN)s, can be prepared by dissolving 
precipitated mercuric oxide in a solution of hydrocyanic acid, 
HCN. It decomposes into mercury and cyanogen, C2Na, 
when heated. 


MAGNESIUM, ZINC, CADMIUM AND MERCURY 491 


* Mercuric Fulminate, Hg(ONC)s, is used in cartridges and 
percussion caps for firearms and in detonating caps for firing 
dynamite and nitroglycerin. 

It is hydrolyzed by hydrochloric acid and water to hydroxyl- 
amine hydrochloride, NH2OH.HCI, and formic acid, HCO2H. 


O—N=C 
He +4HCl+ 44,0 
O—N=C 
ou 
= HgCh + 2H—O—NH, HCI + 20=C¢ 
H 


Ionization of Compounds of Cadmium and Mercury. For 
some reason, not understood, the chlorides and sulfates of cad- 
mium and mercuric mercury ionize to a much smaller degree 
than the corresponding salts of most other metals. 

Solubility of the Sulfides of Group II. The sulfides of the 
metals of the first division of Group II, CaS, SrS and BaS, are 
hydrolyzed by water, forming hydroxides and soluble hydro- 


sulfides : 
2 CaS + 2 HOH = Ca(OH). -++ Ca(SH)o 


Magnesium sulfide, MgS, gives with water magnesium hy- 
droxide, Mg(OH)s, and hydrogen sulfide. Zinc sulfide, ZnS, is 
not affected by water, but dissolves in strong acids, if not too 
dilute. It is almost insoluble in such a weak acid as acetic acid. 
Cadmium sulfide, CdS, dissolves readily, especially on warming, in 
moderately concentrated, strong acids, especially in nitric acid 
(5 per cent) or sulfuric acid (15 per cent). Mercuric sulfide, HgS 
does not dissolve, even in boiling nitric acid, but dissolves 
easily in aqua regia. These relations furnish a ready means of 
separating magnesium, zinc, cadmium and mercury from each 
other and cause them to be classified in three different groups 
for analytical purposes. 

Conduct of Solutions of Magnesium, Zinc and Cadmium 
Salts toward Ammonium Hydroxide. Ammonium hydroxide 
gives no precipitate with salts of these metals in solutions con- 


492 A TEXT BOOK OF CHEMISTRY 


taining ammonium chloride. The zinc and cadmium salts form 
complex compounds, which are soluble, such as Zn(NHs3)4SO,4 
and Cd(NHs3)sCle. These resemble the corresponding com- 
pounds of copper but are colorless. For Mg see p. 479. 

Ammono-mercuric Compounds. When a solution of an alka- 
line hydroxide is added to a solution of a mercurous or mercuric 
salt, mercurous oxide, Hg.O, or mercuric oxide, HgO, is precipi- 
tated, as has been stated. If ammonium hydroxide, NH,OH, is 
added to such a solution, however, compounds of a wholly differ- 
ent type, called ammonobasic mercuric compounds, are precipi- 
tated. These may be considered as formed by the ammonolysis 
of mercuric salts by a process which is closely analogous to the 
formation of an ordinary basic (aquobasic) salt by hydrolysis. 
Thus the partial hydrolysis of mercuric chloride may give in 
solution : 


HgCl, + H.OH = H—O—HgCl + HCl 
Aquobasic 
Mercuric Chloride 


In the presence of ammonia, by an exactly analogous reaction, 
we should have: 
HgCl, + H.NH, = H2N—HegCl + HCl 


Ammonobasic 
Mercurie Chloride 


or Hg(NOs)2 + H.NH, =~ H,N—HgNO; + HNO; 
Ammonobasic 
Mercuric Nitrate 


The hydrochloric or nitric acid would, of course, unite with 
the excess of ammonia present to form ammonium chloride, 
NH.Cl, or ammonium nitrate, NHs.NO;. Salts of many other 
metals undergo ammonolysis in solutions in anhydrous ammo- 
nia, but the ammonobasic compounds which are formed are 
decomposed by water in almost all cases, while the ammono- 
basic mercuric compounds are stable in the presence of water, 
elther because of their extreme insolubility or because of some 
specific affinity between mercury and nitrogen. (See E. C. 
Franklin, J. Am. Chem. Soc. 29, 35 (1907); Am. Chem. J. 47, 
363 (1912) ). 


MAGNESIUM, ZINC, CADMIUM AND MERCURY 493 


Mercurous salts react with ammonia as though they were 
mixtures of a mercuric salt with mercury : 


Hg.Cl, + 2 H.NH, = H.N.HeCl + Hg + NH,Cl. 


The metallic mercury colors the precipitate formed from 
mercurous salts black. 

If ammonobasic mercuric chloride is dissolved in a solution 
of ammonium chloride in anhydrous ammonia, the ammonolysis 
may be reversed exactly as the hydrolysis of a salt may be 
reversed by hydrochloric acid: 


H:N—Hg—Cl + NH,Cl == HgCl, + 2 NH3 


From such a solution a compound of the formula HgCle.2 NHs3, 
which contains ammonia of crystallization and is closely analo- 
gous to the hydrates of other salts (p. 82), may be crystallized. 

Nessler’s Reagent. Mercuric iodide, HglIs, which is almost 
wholly insoluble in water, dissolves easily in a solution of potas- 
sium iodide, owing to the formation of a complex salt, potassium 
mercuric iodide, KeHgI,. In such a solution sodium hydroxide 
will give no precipitate ; but if ammonia or an ammonium salt is 
added to the alkaline solution, a precipitate of ammonobasic- 
aquobasic-mercuric iodide, HO—Hg—NH—Hg—I, is formed. 
In very dilute solutions of ammonia the solution, which is called 
““Nessler’s reagent,” produces a brown coloration which is used 
for the detection and quantitative estimation of ammonia. 


EXERCISES 


1. How much crystallized hydrate of magnesium chloride and how 
much ammonium chloride will be required to furnish one pound (453 
grams) of anhydrous magnesium chloride ? 

2. How many liters of carbon dioxide at 20° and 760 mm. will be 
given by heating 84 grams of magnesium carbonate ? 


3. How much dolomite would be required to give a pound of Epsom 
salts ? 

4, What volume of gases will 0.284 gram of mercuric fulminate give 
by its explosion, supposing the temperature of the gases to be 546°? 


CHAPTER XXVIII 


METALS OF GROUP III. ALUMINIUM FAMILY. RARE 
EARTH METALS 


WHILE all of the elements of both divisions of Group II are fully 
metallic in character and all except radium are comparatively 
common and their compounds well known, the first element, 
boron, of Group III, is decidedly nonmetallic, and aluminium is the 
only metal of the group which can be considered very common. 

Aluminium, Al, 26,97, is found in a great variety of natural 


silicates, especially in the feldspar and mica of the granites and _ 


similar rocks, which are still abundant and which must have 
been much more common in early geologic time. By the pro- 
longed action of water and the forces of nature such rocks have 
been slowly disintegrated. The potassium and sodium of the 
minerals, have been partly, though by no means completely, 
dissolved and removed, and a hydrated silicate of aluminium, 
mixed with fragments of quartz and of partially decomposed 
minerals, has been left in an extremely fine state of division. 
This material, after transportation for some distance by water, 
has been deposited by sedimentation and has formed immense 
beds of shales, clays and soils. Such clays and shales may be 
considered as ores of aluminium, though they contain only from 
15 to 30 per cent of the metal. Pure kaolin, the mineral basis of 
clay, has the formula Al2Si,07.2H:O. Aluminium also occurs as 
the oxide, Al,O3, in crystals known as ruby and sapphire, which 
are used as gems, and in a massive, very hard form, called emery 
and used as an abrasive. The mineralogical name of the oxide 
is corundum. Bauxite, the hydrate, Al.O3.2 H.O, usually 
containing a considerable amount of the hydrate of ferric oxide, 


2 Fe:O3.3 HO, is the chief source from which aluminium oxide 
494 


——— 


ALUMINIUM 495 


is prepared for the manufacture of the metal. Cryolite, Na3AlFs, 
is a soft, easily fusible mineral found in large quantities in Green- 
land, but, so far as known, nowhere else, except as a rare mineral. 

Metallurgy. Aluminium was first prepared by the German 
chemist Wohler in 1828 by the action of potassium on aluminium 
chloride, but he obtained only a very small quantity in the form 
of a gray powder. ‘Twenty-six years later Sainte-Claire-Deville 
exhibited in Paris a quantity of the metal, which he obtained by 
the action of sodium on the chloride, and the element aroused a 
great deal of interest as “ silver from clay.’ The desire of ob- 
taining the metal in larger quantities led to the development of 
cheaper and better methods for the manufacture of sodium, but 
as the valence of sodium is one while that of aluminium is three, 
and the atomic weights are not far different, it must always take 
about three pounds of sodium to give one pound of aluminium, 
and the metal manufactured by that process was never put on 
the market at a price below $10 to $12 a pound. 

In 1885 Professor Mabery of Cleveland, at a meeting of the 
American Association for the Advancement of Science held in 
Ann Arbor, gave an account of a new electric furnace devised 
in 1882, by the Cowles brothers, for the production of aluminium 
bronze. They had discovered that, at the high temperature 
of the electric arc, aluminium oxide, Al,O3, can be reduced by 
carbon to the metallic form, and that if copper is present the alloy, 
aluminium bronze, can be obtained. This seems to have been 
the first application of the electric furnace to an industrial pro- 
cess. Its use for this particular purpose was short-lived. A 
few years later another American, C. M. Hall, discovered that 
aluminium oxide dissolves easily in melted cryolite, and that if 
an electric current is passed from a carbon anode through the 
molten mass contained in an iron pot, aluminium is deposited 
in the bottom of the pot, while oxygen liberated at the anode 
combines with the carbon and escapes as carbon dioxide (Fig. 
103). Other materials are now used wholly or in part in place 
of the eryolite, but the principles used in the process are not 
changed. The aluminium oxide is obtained by heating bauxite 


496 A TEXTBOOK OF CHEMISTRY 


with carbon in an electric furnace. The iron, silicon and other 
elements in the bauxite are reduced by this process and may 
be separated from the 
fused, pure aluminium 
oxide. 

* Thus far the alu- 
minium oxide used for 
the production of alu- 
minium has not been 
prepared, commercially, 
from clay, but has 
usually been made from 
bauxite. Very recently 
NMR eas Nita Oe Alfred H. Cowles has 

Fig. 103 developed a process by 

which clay mixed with 

salt and charcoal is heated in a current of air and steam, giving 
the reaction : 


Al.Si207 + 4 NaCl + 2 H,0O = 2 Na.O.2 SiO2.Al,03 + 4 HCl 


The carbon burns to carbon monoxide and serves to render the 
material porous and easily accessible to the steam and air. Iron, 
which is present, volatilizes as ferric chloride, FeCls. 

If the mixture of sodium silicate and aluminate is mixed with 
lime and heated, an insoluble calcium silicate and soluble sodium 
aluminate are formed : 


2 Na2O.2 SiO»2.Al,O; + 4 CaO 
= 2 Ca,siO4 + NaAlO, + NasAlO3 


Calcium 
Silicate 





Sodium Aluminate 


From the solution of sodium aluminate, aluminium hydroxide 
may be precipitated by carbon dioxide: 
2 NaAlO,. -+- CO. -+ 3 H,O a Al(OH)3 a NaeCO; 


The process seems promising because it gives three valuable 
compounds, hydrochloric acid, sodium carbonate and aluminium 


ew 


ALUMINIUM: THERMITE 497 


hydroxide, with the use of cheap raw materials (Journal of 
Industrial and Engineering Chemistry, 5, 331). 

Properties of Aluminium. Aluminium melts at 658.7° and 
boils at 1800°. It has a specific gravity of only 2.6, almost the 
same as that of glass and scarcely more than one third that of 
iron. This makes it useful for the construction of apparatus 
which should be light. It does not tarnish readily and is used 
to some extent for cooking utensils. It is not attacked by water, 
even at the boiling point, but dissolves readily in alkalies or in 
acids, forming aluminates with the alkalies, such as NaAlOs, and 
salts with acids, such as Al,(SO4)3. Aluminium is rather easily 
corroded by salt solutions. Aluminium which has been amalga- 
mated by bringing it into contact with a dilute solution of mer- 
curic chloride becomes active and will decompose water rapidly 
at ordinary temperatures. In this condition it is in very sharp 
contrast with amalgamated zinc, which is not attacked by hy- 
drochloric or sulfuric acid because of its homogeneous surface. 
The surface of amalgamated aluminium is gray and evidently 
nonhomogeneous. Aluminium is used to some extent for elec- 
tric conductors in place of copper. It is often used as an addi- 
tion to cast iron, greatly improving its quality. 

Alloys. The best known alloy is aluminiwm bronze, composed 
of copper with 5-12 per cent of aluminium. It resembles gold 
very closely in appearance and does not tarnish readily. Magna- 
lium, an alloy with a small amount of magnesium, is very light 
and is much more easily worked on a lathe than aluminium 
itself. Alloys containing from 2 to 10 per cent of copper are 
used for castings for automobiles and for other purposes where 
lightness is desirable. 

Goldschmidt’s Thermite Process. Aluminium has a very 
strong affinity for oxygen, as shown by the difficulty with which 
it is reduced. The heat of combustion of aluminium is: 


2 Al+ 30 = AI.03 + 380,000 calories 
That of iron, if it could be burned to ferric oxide, is: 
2Fe + 3.0 = Fe,03 + 195,000 calories 


498 A TEXTBOOK OF CHEMISTRY 


From these values it is evident that the reaction 
Fe.O3 + 2 Al = Al,O3 oe Z Fe 


may occur with the evolution of a large amount of heat. Gold- 
schmidt has made use of this principle for the production of very 
high temperatures and also for the reduction of chromic oxide 
and other refractory oxides. Because the aluminium oxide is 
not volatile, the heat of the reaction is not dissipated by the for- 
mation of a vapor, and a temperature high enough to melt iron 
or steel may be easily obtained for the welding of steel rails, per- 
foration of iron plates and similar purposes. 

The thermite process is also very useful for the production of 
chromium and other metals which it is difficult to obtain in 
other ways. A special advantage of the process, in some cases, 
is that the metals obtained in this way are free from carbon. 

Aluminium Chloride, AICl3, is easily prepared by heating 
aluminium turnings in chlorine or in hydrochloric acid. It 
sublimes at 183° under atmospheric pressure and melts at 193° 
under increased pressure. Aluminium chloride dissolves in 
water with the evolution of considerable heat. If the solution 
is not too dilute, the addition of concentrated hydrochloric acid 
will cause a crystalline hydrate, AlCl;.H.O, to separate. On 
heating this hydrate it loses hydrochloric acid and aluminium 
oxide is left: 


Aluminium chloride forms with many organic substances 
extremely reactive addition compounds which are used in a 
variety of syntheses. For these reactions water must be care- 
fully excluded, and the hydrate of the chloride cannot be used 
at all for such purposes. All of these facts indicate very clearly 
that the water of the hydrate is in a state of intimate chemical 
combination such as to greatly modify the relation which the 
chlorine and aluminium have in the anhydrous chloride. Solu- 
tions of aluminium chloride are partially hydrolyzed by water, 
and react strongly acid. Nearly all salts of trivalent or quadri- 
valent metals act in the same way. 


ALUMINIUM COMPOUNDS 499 


Aluminium Fluoride, AlF3. Either aluminium or aluminium 
hydroxide dissolves easily in an aqueous solution of hydrofluoric 
acid, forming a supersaturated solution from which anhydrous 
aluminium fluoride slowly separates in small crystals. The 
double fluoride, Na3AlFs, is found in nature as the mineral 
eryolite, and was formerly used to dissolve the oxide for the 
electrolytic preparation of metallic aluminium. 

Aluminium Hydroxide, Al(OH)s, is precipitated from solutions 
of aluminium salts on the addition of an alkaline hydroxide. 
It forms a voluminous, gelatinous precipitate which dissolves 
either in solutions of strong acids or strong bases. For this 
reason it is called amphoteric, meaning that it has both basic and 
acid properties. It seems probable that in contact with an acid 
its hydroxyl combines with the hydrogen of the acid, while in 
contact with a base its hydrogen combines with the hydroxy] of 
the base. The compound may be considered, therefore, either 
as a triacid base or as a tribasic acid. It is, of course, very weak 
both as an acid and asa base. Salts in which it is the cation are 
hydrolyzed in solution and have an acid reaction : 


Al,(SO.)3 + 6 HOH = 2 Al(OH); + 3 H.S0, 


Those salts in which the aluminium forms part of the anion are 
also hydrolyzed and have an alkaline reaction : 


Na;AlO; + 3 HOH = Al(OH); + 3 NaOH 


When heated, aluminium hydroxide loses water and is con- 
verted into the oxide, Al.O3. It conducts itself very much as 
the silicic acids, however, losing a part of the water very easily 
at ordinary temperatures but requiring ignition at bright redness 
to expel the last portions. 

Aluminium Oxide, Al.O3, is found in nature as the mineral 
corundum. In its massive, not very pure forms, it is called 
emery. Next to the diamond it is the hardest mineral known 
and has been long used as an abrasive for grinding and polish- 
ing glass and metals. It has now been partly displaced for these 
-uses by carborundum, SiC, which is much harder. Crystalline 


UNIVERSITY OF ILLINOIS 
PIBRARY-CHEMISTRY 


500 A TEXTBOOK OF CHEMISTRY 


forms of corundum colored blue by some foreign substance are 
called sapphires, or other forms colored red by chromium are 
called rubies. The latter are now made artificially. A fused 
oxide, prepared in the electric furnace, is used, under the name 
of alundum, as an abrasive and also as a refractory material. 

Ignited aluminium oxide is insoluble in acids but may be 
brought into solution slowly by fusion with sodium pyrosulfate, 
Na2S.O07. Aluminium oxide is reduced by carbon at the tem- 
perature of the electric furnace, but cannot be reduced at lower 
temperatures. Before metallic aluminium was prepared by 
electrical processes, the anhydrous aluminium chloride used for 
the preparation of the metal was obtained by heating a mixture 
of the oxide with carbon in a current of chlorine: 


Aluminium Sulfate, Al.(SO,)3.18 H:O, is prepared by the 
decomposition of clay with sulfuric acid. A more or less pure 
sulfate containing some ferric sulfate is extensively used under 
the name of “alum”? for the clarification and purification of 
water. If a solution of the sulfate is mixed with a water con- 
taining calcium bicarbonate, an insoluble, gelatinous precipitate 
of aluminium hydroxide is formed : 


Als(SOx)s3 ++ a CaH2(CQO3)> — Tey Al(OH)3 aa 3 CaSO, +- 6 CO, 


The precipitate collects fine particles of clay and also bacteria 
which are suspended in the water, and by suitable filtration clear 
water, nearly or quite free from disease germs, is obtained. The 
amount of aluminium sulfate required is so small that the calcium 
sulfate formed does not seriously increase the permanent hard- 
ness of the water. All of the aluminium added is removed in 
the filtration. 

Alums, M’M’’’(SO,)..12 H.O. By adding potassium sul- 
fate, K:50O., to a solution of aluminium sulfate a compound 
having the composition KAI(SO4)2.12 H.O, and known since 
early times under the name of alum, is formed. It crystallizes 
in octahedra, which, with care, may be obtained in very perfect 
forms of large size. 


ALUMS, EARTHENWARE 501 


Alum was formerly much used as a mordant in dyeing. The 
aluminium hydroxide formed by its hydrolysis combines with 
many coloring matters to form insoluble compounds called 
lakes. ‘These compounds attach themselves strongly to the 
fibers of the cloth and cannot be removed by washing. Alu- 
minium sulfate and the aluminates have largely displaced alum 
for such uses because the potassium sulfate is expensive and 
unnecessary. 

The potassium of alum may be replaced by ammonium or 
other univalent metals, the aluminium may be replaced by ferric 
iron or other trivalent metals and even the sulfate radical, SOx, 
may be replaced by the selenate radical, SeO,. This gives 
a great variety of alums, all of which crystallize in octahedra 
and are isomorphous. A crystal of any alum will grow in a 
supersaturated solution of any other. 

The following may be given as illustrations of the alums: 


Ammonium Alum NH, Al(SO4)2 . 12 H2O 
Ammonium Ferric Alum NHyFe(SOx)2 . 12 H2O 
Chrome Alum KCr(SO,)2. 12 H2O 
Rubidium Alum RbAI(SO,)2 . 12 H.O 


Brick, Earthenware, Porcelain. Aluminium silicate melts 
only at very high temperatures, but the presence of other com- 
mon metals, such as iron, calcium, magnesium, sodium or potas- 
sium, lowers the melting point. Ordinary clays contain com- 
pounds of these metals distributed as very fine particles through- 
out their mass, and when such clays are heated to a high 
temperature, the particles melt and cause the material to sinter 
together to a strong but very porous mass. In addition to this 
property of sintering without fully melting, the original clays 
become plastic when mixed with water, and in this condition 
may be molded into bricks or into the “ biscuit ”’ forms which 
furnish the basis of earthenware or porcelain. The plasticity 
of the clay seems to be closely connected with its colloidal 
character. 

In order to give them a surface which is smoother and im- 


02 A TEXTBOOK OF CHEMISTRY 


pervious to water, articles of earthenware and porcelain must 
be covered with a glaze. Several methods of glazing are in 
use. One method is to throw salt into the furnace after the 
biscuit has been well burned. The sodium chloride reacts with 
the silica and alumina of the clay and the moisture of the air 
to form hydrochloric acid, which escapes, and a fusible silicate 
is formed, which melts and covers the surface with a glass. 
Other glazes are made from mixtures containing lead oxide. 
Some glazes of this type are not wholly insoluble in water, and 
England has enacted stringent laws requiring lead glazes to be 
highly insoluble when the articles are to be used to contain or 
cook food. Porcelains are usually glazed by the application of 
finely powdered feldspar and subjecting them to a temperature 
which causes it to melt and run into the surface. 

Ultramarine. Small quantities of a beautiful blue stone 
called lapis lazuli are found in nature. When powdered this 
stone gives a beautiful blue pigment which is not affected even 
by long exposure to the light. The mineral is so rare, however, 
that during the first years of the nineteenth century the pig- 
ment was sold to artists at $60 an ounce. In 1828 Gmelin dis- 
covered that the material can be made artificially by heating 
mixtures of clay, sodium sulfate, charcoal and sulfur. The 
artificial product is fully equal to the natural mineral for the 
uses to which it is applied and is now sold at a few cents per 
pound. By changing the method of manufacture, other com- 
pounds, ranging in color from reddish violet to bluish green, 
are also made. In spite of a very large amount of work devoted 
to the preparation:and analysis of these compounds it has not 
been possible to assign definite formulas to them. 

The Rare Earths. These include the oxides of a number of 
elements all of which are characterized by being trivalent and 
by forming oxalates which are insoluble in dilute mineral acids. 
They resemble each other very closely in all their properties 
and in the types of compounds which they form. Their salts 
are isomorphous and do not as a rule differ greatly in solubility. 
Because of these slight differences they cannot, with the possible 


a - 


ALUMINIUM FAMILY. RARE EARTH METALS 6503 


exception of cerium, be separated from each other quantita- 
tively, and their preparation in a pure state requires a long-con- 
tinued series of fractional crystallizations or precipitations. In 
fact, some of them have not been obtained in a high state of 
purity. These elements may be divided into two groups: the 
Cerium group, including cerium, lanthanum, neodymium, pra- 
seodymium, samarium, europium and gadolinium; and the 
Yttrium group, including terbium, dysprosium, holmium, 
yttrium, erbium, thulium, ytterbium, scandium and lutecium. 
The members of the cerium group may be separated roughly 
from those of the yttrium group by making use of the fact that 
the elements of the former form double sulfates with sodium, 
sulfates which are insoluble in a saturated solution of sodium 
sulfate, while the corresponding compounds of the yttrium 
group are soluble. 

* Scandium, Sc, 45.1. When Mendeleeff proposed the Periodic 
System of the elements in 1869, he predicted that several ele- 
ments not then known would probably be discovered in the 
future. Among these was an element which he called “ eka- 
boron” which should have an atomic weight of about 44 and 
form compounds similar to those of aluminium. ‘Ten years 
later Nilson found scandium among the elements found in 
gadolinite and euxenite, and shortly after Mendeleeff pointed 
out that this is in reality the “ ekaboron ” which he had pre- 
dicted. The hydroxide, Sc(OH)3, oxide, Sc203, sulfate, 
Se2(SO.4)3.6 H2O, oxalate, Se2(C20.4)3.6 H2O and other salts 
are known. 

*Yttrium, Y, 89, is found in gadolinite, xenotime and monazite. 
Its compounds resemble those of scandium. The oxide, Y2Os, 
chloride, YCls, sulfide, Y2S3, phosphate, YPOu., and bromate, 
Y(BrO3)3;.9 H2O, may be mentioned. 

* Lanthanum, La, 138.9, is the most positive of the rare earth 
‘metals. Its oxide, La,2O3, combines with water, much as lime 
does, forming the hydroxide, La(OH)s;, which turns litmus paper 
blue. The hydroxide also absorbs carbon dioxide from 
the air, forming the carbonate, Lag(CO3)3. The oxalate, 


504 A TEXTBOOK OF CHEMISTRY 


Lae(C2.04)3.9 HO, is difficultly soluble, as are the oxalates of 
all of the rare earth metals. 

*Ytterbium, Yb, 173.6, gives the sulfate, Ybo(SO.)3.8 H.O, 
carbonate, Ybo(CO3)3.4 HO, acetate, Yb(C2H302)3.4 H2O, oxa- 
late, Yb2(C2O4)3.10 HO, and many other salts. 

* Praseodymium, Pr, 140.92, and Neodymium, Nd, 144.27. 
In 1842 Mosander obtained from cerite (a silicate containing 
cerium, lanthanum, praseodymium and neodymium) an oxide 
of a metal to which he gave the name didymium. In 1885 
Auer. v. Welsbach discovered that by a long series of crystalli- 
zations the double nitrate of the metal previously called didym- 
ium could be separated into two compounds, praseodymium 
ammonium nitrate, Pre(NO3)3.2 NH4NO3.4 H2O, and neodym- 
ium ammonium nitrate, Nd(NO3)3.2 NH4NO3.4H2O. These 
two salts are isomorphous and do not differ very greatly in 
solubility. It was necessary, therefore, to repeat the crystalli- 
zation many hundreds of times in such a manner that the more 
soluble portions were transferred in one direction through the 
crystallizing dishes and the less soluble portions in the other 
direction, while portions of the same degree of separation were 
systematically united. This process of fractional crystalliza- 
tion has been much used in the separation of the metals of the 
rare earths. (See Auer. v. Welsbach, Monatshefte fur Chemie, 
6,477; Baxter and Chapin, J. Amer. Chem. Soc. 33, 5; James, 
ibid. 30, 182; 31, 913.) 

Typical compounds of praseodymium and neodymium are, 
Pr.Os, PrO,, Nd.Os, Pr2(SO,4)3.8 HO, Nd2(SQx.)3.8 H.O, 
Pr2(C2O4)3 ; 10 HO, Ndo(C204)3 : 10 HO, Pr(BrQs3)3 - 9 H,0, 
Nd(BrQs3)3 6 9 H,O. 

The oxalates of nearly all of the rare earth metals are very 
difficultly soluble and are often used as a means of separating 
these metals from those of other groups. 

The salts of praseodymium are green in color, those of: 
neodymium are rose-colored, the two colors being comple- 
mentary very much as those of the salts of cobalt and nickel 
are. 


ALUMINIUM FAMILY. RARE EARTH METAIS 505 


*Tllinium, Il. When the neodymium and samarium com- 
pounds obtained by B.S. Hopkins, L. F. Yntema and J. A. Harris 
after a long series of fractional crystallizations, starting with 
the mixture of sulfates which is a by-product in the manufacture 
of Welsbach mantles, were examined spectroscopically at the 
Bureau of Standards, more than 100 spectral lines common to 
both samples but, apparently, characteristic of neither element, 
were found. Further fractionation gave a wine-colored solu- 
tion having a characteristic absorption band not far from the 
position of the sodium line in the spectrum. An examination 
of the X-ray spectrum of the oxide of the metal in this solution 
gave two X-ray spectral lines in the positions which theory 
predicts for the element of atomic number 61, the only missing 
element in the rare earth series. Compounds of the element 
have not yet been sufficiently purified to admit of a determina- 
tion of the atomic weight. This should be approximately 146. 

* Samarium, Sm, 150.4, is found in samarskite, a columbate of 
metals of the rare earths, and was first partially separated from 
the “ didymium ” of that mineral by Lecoq de Boisbaudran. 
It is less basic than praseodymium and neodymium, and its 
double nitrate with magnesium is more easily decomposed by 
heat, a method sometimes used in separations. The oxide, 
Sm,.Os, and solutions of its salts are yellow. It forms a chloride, 
SmCl:, in which the metal is bivalent, but in nearly all of its 
salts it is trivalent. A considerable number of salts have 
been prepared, such as samarium sulfate, Sme(SOx)3.8 HO, 
the nitrate, Sm(NOs3)36 HO and the carbonate, 
Sm2(CQO3)3.3 HO. 

*Europium, Eu, 152, Gadolinium, Gd, 157.3, and Terbium, 
Tb, 159.2, form a group of weakly basic earths intermediate 
between the cerium earths on the one side and the yttrium 
earths on the other. They have been separated by tedious 
fractional crystallizations and finally identified by means of 
their spectra and determinations of their atomic weights. 
Europium oxide, Eu,O3, has a light rose color, and the sulfate, 
Eu2(SOx4)3.8 HO, is also rose-colored. Gadolinium oxide, 


506 A TEXTBOOK OF CHEMISTRY 


Gd.O3, and sulfate, Gdo(SO.4)3.8 HeO,. are white. Terbium 
oxide, Tb2O3, is also white; but a higher oxide, possibly TbOkz, 
but not yet obtained pure, is dark brown or black according to 
the method of preparation. 

* Holmium, Ho, 163.4, is obtained from euxenite, a columbate 
and titanate of the yttrium and cerium earths. The solubility 
of the double sulfate with ammonium seems to lie between 
those of yttrium and erbium (Holmberg). Apparently pure 
compounds have not yet been prepared. The metal is named 
from Stockholm. 

*Dysprosium, Dy, 162.5, has been obtained in its purest 
form by fractional crystallization of salts of ethyl sulfuric acid, 
H(C2H;)SOy. The dysprosium salt is Dy(Cs:H;SO,)3. The 
oxide, Dy.O3, is white. The bromate, Dy(BrOs;)3.9 HO, is 
yellow. 

* Erbium, Er, 167.7, is found among the yttrium earths from 
euxenite and other sources. The oxide, Er.Q3, is rose-colored, 
as are also the sulfate, Ero(SO.)3.8 H2O, and other salts. 

*Thulium, Tu, 169.4, is found in euxenite and other rare 
earth minerals. It seems to have been separated in the purest 
condition by the fractional crystallization of the bromates 
(James, J. Amer. Chem. Soc., 33, 13832). The oxide, TuyQs, is 
white, with a faint green tint. The bromate, Tu(BrO3)3.9 HO, 
separates in pale, bluish green prisms, isomorphous with the 
bromates of other rare earth metals. 

* Lutecium, Lu, 175, is one of the more recently discovered 
elements of this group and has been separated from the gado- 
linite earths. 

* Gallium, Ga, 69.72, was also predicted by Mendeleeff under 
the name of “ eka-aluminium.”’ Unlike the other metals of the 
group thus far described, it forms two classes of compounds, 
those in which it is bivalent and others in which it is trivalent. 
The chlorides are GaCl, and GaCl3; the sulfates, GaSO, and 
Gap(SOx)3.18 HeO. The former is oxidized by potassium perman- 
ganate as ferrous sulfate is. The alum, NH4Ga(SO,)o.12 HO, 
is isomorphous with ordinary alum. 


ALUMINIUM FAMILY. RARE EARTH METALS 507 


Indium, In, 114.8, was discovered by means of the blue line 
of its spectrum by Reid and Richter shortly after the methods 
of spectrum analysis had been developed by Bunsenand Kirchoff. 
The analysis of its oxide gave about 76 parts of indium for 16 
parts of oxygen, and an atomic weight of 76 was at first assigned 
to the element. But Mendeleeff pointed out that this would 
place it between arsenic and selenium, where there is no vacant 
place in the Periodic System, and also that its properties did not 
agree with such a position in the table. He suggested, therefore, 
that the formula of the oxide is IngO3 and the atomic weight 
114. The determination of the specific heat gave the value 
0.056. This points to an atomic weight of 6.2/0.056 = 110, 
which agrees fairly well with an atomic weight of 114 but would 
not agree at all with the value of 76. The preparation of an 
ammonium alum, NH4In(SO,4)o.12 H.O, soon after this, gave 
further support for the accepted formula. 

Indium is a soft, white metal, which melts at 155°. It gives 
three chlorides, InCl, InCh, InCl;, but the compounds in which 
it is trivalent are most stable and best known. 

Thallium, Tl, 204.39, was also discovered by means of the 
spectroscope. Crookes found it in 1861 in the slimes from sul- 
furic acid made at Tilkerode in the Harz. He named it thallium 
from the Latin word thallus, meaning a young twig, because of 
a brilliant green line in its spectrum. Metallic thallium is a 
bluish white, soft metal, somewhat resembling lead. 

Thallium forms thallous compounds in which it is univalent, 
and thallic compounds in which it is trivalent. Of the former, 
thallous oxide, Tl,O, thallous chloride, TICI, thallous hydroxide, 
TIOH.H.O, and thallous sulfide, ThS, may be mentioned. 
The last is a black precipitate nearly insoluble in acetic acid 
but soluble in mineral acids. Thallous iodide, TII, is also very 
difficultly soluble. 

Among the thallic compounds are thallic chloride, TICl;.H.O, 
thallic nitrate, T1(NO3)3.8 HO, and thallic sulfide, TlSs. 


CHAPTER XXIX 


TIN AND LEAD 


Ir has been pointed out (p. 361) that tin and lead belong to 
the carbon group of the Periodic System and that each gives 
an oxide (SnOz and PbO) resembling carbon dioxide, COs, and 
silicon dioxide, SiOz, in formula and in some other properties, 
especially in their acidic character. Both of these elements 
are clearly metals rather than non-metals in most of their 
properties. 

Tin (Sn, 118.70). Occurrence, Metallurgy. Tin is rather 
remarkable in that, although it is an element which is found in 
sufficient quantity so that it is a common metal for household 
and commercial use, there are only a very few localities in the 
world where its ores can be profitably mined. One of the oldest 
of these is Cornwall in England, where tin has been obtained 
for nearly or quite twenty centuries and which furnished a large 
part of the tin used in the world until comparatively modern 
times. ‘The world’s supply of tin for one year is approximately 
80,000 tons, and this comes almost entirely from Banca and the 
East India islands, Tasmania, Bolivia and Cornwall. The only 
important ore is cassiterite, stannic oxide, SnOz. This is found 
sometimes in veins, sometimes as a heavy gravel, called “ stream 
tin.” The metallurgy is comparatively simple, consisting in the 
reduction of the oxide by means of charcoal or coal. On ac- 
count of its value, the recovery of tin from tin scrap has also 
assumed considerable commercial importance. Several methods 
are in use, one of the best being the treatment of the scrap with 
dry chlorine gas, which converts the tin into stannic chloride, 
SnCl4, but leaves the iron comparatively unattacked. Stannic 


chloride is volatile and can be easily separated. 
508 


TIN 509 


Uses of Tin. Alloys. Tin Plate. Tin is not affected by dry 
or moist air or by water, even at the boiling point or higher. 
It is, for this reason, the most suitable of the cheaper metals 
for the tubes of condensers to be used in the preparation of 
distilled water. The principal use of the metal is in the manu- 
facture of tin plate — sheet iron which has been covered with a 
thin coating by dipping the carefully cleaned metal in a vat of 
melted tin. Tin is more electropositive than iron, hence when 
the two metals are in contact with water or a dilute acid the 
tendency is for the iron to corrode while the tin is protected. 
For this reason tin vessels rust through rapidly as soon as the 
iron is exposed at any point — exactly contrary to the conduct 
of iron coated with zine (p. 482). ‘‘ Terne plate,’ which is 
used for roofing purposes, is covered with an alloy of lead: and 
tin, the lead being used because it is very much cheaper than tin. 

Solder is an alloy of lead and tin, used to join pieces of tin 
plate in making culinary vessels of all kinds. Common solder 
contains equal parts of the metals, but fine solder, containing 
more tin, and coarse solder, containing more lead, are often used. 

Tin was formerly much used in gun metal and bell metal and 
is still used in statuary and ornamental bronzes, which are alloys 
of tin and copper, usually with a little lead and zine. 

Tin, antimony and lead are the principal constituents of 
Britannia metal, pewter and Babbitt metal. Copper and other 
metals are sometimes added. 

Tin melts at 231.9°. It is convenient to remember the melting 
points of tin (232°), lead (327°) and zine (419°) as about 100° 
apart, with tin the lowest and zine the highest. Tin boils at a 
high temperature, but the boiling point has not been determined. 
The specific gravity is 7.30. 

The metal oxidizes slowly, when heated above its melting 
point in the air, to stannic oxide, SnO». It dissolves as stannous 
chloride, SnCl:, in concentrated hydrochloric acid and is con- 
verted by nitric acid into a mixture of stannic and metastannic 
acids, SnOe.H2O, which is insoluble in an excess of the acid or 
in water. 


510 A TEXTBOOK OF CHEMISTRY 


Compounds of Tin. Tin forms stannous compounds, such as 
SnCly, in which it is bivalent, and stannic compounds, as SnClu, 
in which it is quadrivalent. In the former it is rather strongly 
basic and metallic, in the latter much weaker as a base and in 
some of the compounds distinctly acidic. 

* Stannous Oxide, SnO. If a solution of potassium carbonate, 
K,CO3, is added to a solution of stannous chloride, SnCle, a 
white precipitate having the composition 2 SnO.H.2O is formed ; 
but this loses water on heating the solution, especially if a 
little alkali is present, much as cupric hydroxide, Cu(OH)s, does 
(p. 431), and is changed to black insoluble stannous oxide, SnO. 
The original precipitate is amphoteric, like aluminium hydroxide, 
Al(OH)3, and dissolves either in acids or in alkalies. 

Stannous Chloride, SnCle.2 H2O, is easily obtained by dis- 
solving tin in concentrated hydrochloric acid. The anhydrous 
salt can be prepared by heating this hydrate in a stream of 
hydrochloric acid. It boils at 606°. 

Stannous chloride is frequently used in the laboratory as a 
reducing agent because of its strong tendency to take up chlorine 
or other.elements and pass over to the stannic form. A solution 
of stannous chloride to which an excess of sodium hydroxide 
has been added, forming sodium stannite, NaHSnOz or NazSnOz, 
is also a powerful reducing agent. 

Stannous chloride and several other compounds of tin are ex- 
tensively used as mordants in dyeing. 

Stannous Sulphide, SnS, is precipitated from acid solutions of 
stannous salts as a dark brown, almost black compound It 
does not dissolve in colorless ammonium sulfide, (NH4).S, but 
the yellow ammonium or sodium sulfides, which contain poly- 
sulfides, (NH4)2Ss, ete., dissolve it as ammonium or sodium 
sulfostannate, (NH4)2SnS3 or NaoSnS3. From such a solution 
acids precipitate yellow stannic sulfide, SnSy. 

Stannic Oxide, SnOz, is prepared by the oxidation of tin in 
air at a high temperature or by treating tin with nitric acid and 
igniting the mixture of stannic acids which is formed. It is 
also found as the crystalline mineral, cassiterite, in nature. 


STANNIC ACIDS oll 


Stannic oxide does not dissolve in the melted silicates which 
form glass, and it has been sometimes used for the manufacture 
of white, opaque glass, but less expensive materials are usually 
employed. 

Stannic Acids. Tin resembles silicon in that several acids 
are derived from the same anhydride, SnOse, just as there is a 
long list of silicic acids derived from silicon dioxide, SiQs. 
Berzelius discovered in 1817 that the stannic acid, obtained by 
precipitation from a solution of stannic chloride, SnCly, is very 
different in its properties from metastannic acid, which is formed 
by treating tin with nitric acid, and his study of these compounds 
led him to propose the word zsomer to designate a compound 
having the same composition as some other compound which 
has different properties. At that time the compounds which 
we now call anhydrides were called acids, and the compounds 
which he considered isomeric were the anhydrides of the two 
acids rather than the compounds which we should now call 
stannic and metastannic acids. Later investigations have 
shown that neither the free acids, when dried in the air, nor 
their salts are isomeric as the term is used to-day. This will be 
clear from the following table. A third acid, parastannic acid, 
which was also discovered by Berzelius, is included. Stannic 
and metastannic acid are isomeric when dried in a vacuum. 








CHLORIDE 





Nt hain FORMULA DRIED | FORMULA DRIED FORMULA OF Ta Ruaare 
IN THE AIR IN A VACUUM POTASSIUM SALT HCl 

pt ee Es ee ee ae | éC [ae 
Stannic acid | H2eSnO3.H2O H2Sn0O3 K2Sn0O3.H20 SnCla 
Metastannic 

acid. . . | HeSnsOnu.9 H2O | H2SnsOu 4 H2O | KeSnsOn.4 H2O SnsOeCle .4H20 
Parastannic 

acid. . .| HeSnsOn.7 H2O | HeSnsOu.2 H20 | KeSnsO1.20r3 H2O | SnsOcCle.2 H2O 





Stannic Acid, H,SnO3.H2O, or H.SnO,, is obtained by pre- 
cipitating a solution of stannic chloride, SnCly, with ammonia 
or with calcium carbonate. It dissolves easily in strong acids 
or in alkalies. From its solution in alkalies it is reprecipitated 


512 A TEXTBOOK OF CHEMISTRY 


by acids. On drying it is partly changed to metastannic acid, 
and a failure to understand this has led to much confusion in 
the literature. 

Metastannic Acid, H2Sn;01;.9 H2O or (H4SnOx)s, is the prin- 
cipal product formed by the action of warm nitric acid on tin. 
If the mixture obtained in this way is dissolved in a little sodium 
hydroxide, the addition of an excess of the alkali will cause the 
precipitation of sodium metastannate, while the sodium stannate 
will remain in solution. 

Metastannic acid is insoluble in nitric acid or sulfuric acid. 
When treated with concentrated hydrochloric acid it forms a 
chloride, Sn;O gCle.4 HO, which dissolves in water but is re- 
precipitated by concentrated hydrochloric acid. Solutions of 
stannic chloride, SnCl4, which have stood for some time, con- 
tain this compound, formed by hydrolysis and rearrangement 
or condensation. It is properly named ‘metastannyl chloride. 

*Parastannic Acid, H.Sn;O0,;.7 H.O, was obtained by Ber- 
zelius by heating metastannic acid with water at 100°. It is 
quite similar to metastannic acid. 

Stannic Chloride, SnCl., is a volatile liquid which boils at 
114° and fumes strongly in the air, owing to its hydrolysis by 
the moisture of the air and the escape of hydrochloric acid. It. 
dissolves in water and forms several hydrates, but it seems pretty 
certain that these contain compounds of the same general char- 
acter as metastannyl chloride (see above) rather than hydrates 
of stannic chloride, as that would ordinarily be understood. 
But our knowledg@ of the structure of hydrates in general is 
still very imperfect. 

Stannic Sulfide, SnSe, separates as a yellow, amorphous pre- 
cipitate when hydrogen sulfide is passed into an acid solution 
of a stannic salt. It dissolves in concentrated hydrochloric 
acid, resembling antimony and differing from arsenic in this 
respect. It also is not precipitated from and dissolves in solu- 
tions of oxalic acid from which antimony can be precipitated as 
the sulfide, SbeS3. Stannic sulfide dissolves in ammonium sulfide 
as ammonium sulfostannate, (NH4).SnS3. 


FIREPROOFING COTTON GOODS 513 


Fireproofing of Cotton Goods. Many fatal accidents occur 
every year from the burning of clothing, and serious accidents 
have occurred from the burning of curtains and fabrics in theaters. 
A variety of substances have been used. to render fabrics less 
inflammable. One of the best of these is stannic oxide. Pro- 
fessor Perkin describes its application as follows: ‘“‘ The flan- 
nelette (or other material) is run through a solution of sodium 
stannate of approximately 45° Tw (sp. gr. 1.225) in such a man- 
ner that it becomes thoroughly impregnated. It is then 
squeezed to remove the excess of stannate solution, passed over 
heated copper drums to thoroughly dry it, after which it is 
run through a solution of ammonium sulfate of about 15° Tw, 
and again squeezed and dried. Apart from the precipitated 
stannic oxide, the material now contains sodium sulfate and this 
is removed by passage through water; the material is then 
dried and subjected to the ordinary processes of finishing. A 
long series of trials, carried out under the most stringent condi- 
tions, have conclusively proved that material, subjected to this 
process, is permanently fireproof. No amount of washing 
with hot soap and water will remove the fireproofing agent, 
or in other words, the property of resisting flame lasts so 
_long as the material itself lasts.” (Address before the Inter- 
national Congress of Applied Chemistry, New York City, Sep- 
tember, 1912.) 

Lead, Pb, 207.2. Occurrence, Metallurgy. Lead is most 
often found as the sulfide, PbS, in the form of galena, a heavy 
black mineral which crystallizes in cubes having a bright, metallic 
luster. Galena is usually associated with other minerals, es- 
pecially with sphalerite, ZnS, and pyrites, FeS:, and with a 
gangue of quartz, SiOz, fluorite, CaF2, barite, BaSOu,, or calcite, 
CaCOQs. | 

Lead is obtained from the ore by roasting it to convert a part 
of the sulfide to the oxide or sulfate : 


PbS + 30 = PbO + SO2 
PbS + 40 = PbsSO, 


514 A TEXTBOOK OF CHEMISTRY 


If the mixture of sulfide with oxide or sulfate is heated with 
exclusion of air or in a reducing atmosphere, the compounds 
mutually reduce each other and sulfur dioxide escapes: 


2 PbO + PbS = 3 Pb + SO, 
PbSO, + PbS = 2 Pb + 2 SO; 


This process is most suitable for very pure ores. Less pure ores, 
which are often reduced for the silver and other metals which 
they contain rather than for the lead, are usually reduced in a 
blast furnace (p. 541) by the combined action of coal or coke 
and iron. The iron combines with the sulfur of the galena, 
reducing it to metallic lead : 


PbS + Fe = FeS + Pb 


The recovery of silver, which is usually present in crude lead, 
has been discussed in a previous chapter. 

Properties and Uses of Lead. Alloys. Lead is the heaviest 
of the cheaper metals, having a specific gravity of 11.34, which 
is even higher than that of silver. It is because of this property, 
and also because of the ease with which lead can be melted and 
cast, that it is used for bullets and shot. Lead melts at 327.4°. 
It has a bright white luster when freshly cut, but tarnishes 
quickly in the air. Lead dissolves easily in nitric acid but only 
slowly and to a slight extent in hydrochloric acid. Even hot 
sulfuric acid scarcely attacks the metal till it has the specific 
gravity of 1.72. A more concentrated acid dissolves lead sulfate 
and attacks metallic lead strongly. 

Lead is so soft that it can be pressed through a die into the 
form of tubing, by means of hydraulic pressure. Such tubing 
is used for waste pipes from sinks and for similar purposes and 
is liked by plumbers because of the ease with which it can be 
worked. It isnot suitable for pipes to convey drinking water, 
because a little of the lead is liable to dissolve, and all soluble 
lead compounds are very poisonous. Lead has the further, 
very dangerous, property, that it acts as a cumulative poison so 
that minute quantities taken daily for some weeks or months 
may finally produce fatal results. 


LEAD nie 


Lead has such slight tenacity that it cannot be drawn into 
wire, in spite of its softness. It can be rolled into sheets and 
beaten into thin foil. Foil made from the alloy with tin has 
sometimes been substituted for pure tin foil to wrap around 
articles of food, but such use is strongly condemned because of 
the poisonous character of the lead. 

The more important alloys of lead, solder, Britannia metal, 
pewter, Babbitt and other antifriction metals, type metal, 
stereotype metal and fusible alloys used for safety fuses, have 
been mentioned in previous chapters. 

Oxides of Lead. ‘There are three definite, well-characterized 
oxides of lead: lead monoxide, or litharge, PbO, lead plumbate, 
usually called red lead or miniwm, Pb3Q0.4, and lead dioxide,' or 
plumbic anhydride, PbO». Three other oxides, Pb2O, Pb;O7 and 
Pb2O3, have been described by various authors, but there is 
considerable doubt whether these represent definite compounds 
or not. The evidence in favor of the existence of the oxide, 
Pb.Os3 (or PbPbOs, lead metaplumbate), is better than that for 
the other two. 

Lead Monoxide or Litharge, PbO, is readily formed by exposing 
lead at a red heat to the action of the air. At that tempera- 
ture the melted film of litharge constantly flows to the side, ex- 
posing a fresh surface to oxidation. The melted litharge is 
collected, allowed to solidify, and is then ground to a fine powder 
for the market. It forms a buff-colored powder used in the manu- 
facture of “boiled ”’ linseed oil and flint glass, and for other 
purposes. 

For the manufacture of red lead, metallic lead is first oxidized 
at a lower temperature, such that the oxide formed does not 
melt. This oxide is then heated to dull redness, but below the 
melting point of litharge, with free access of air. 

The chemical character of red lead is most clearly shown by 
the action of dilute nitric acid upon it. This dissolves two 


1 Often called lead peroxide. It is better to restrict the designa- 
tion peroxide to compounds having a structure similar to that of 
hydrogen peroxide, H—O—O—H. 


516 A TEXTBOOK OF CHEMISTRY 


thirds of the lead and leaves a dark brown soanhes of lead di- 
oxide, or plumbic anhydride, PbO: : 


This reaction shows that two thirds of the lead is basic and 
one third acidic in character, or, in other words, that red lead 
is a lead salt of plumbic acid, H4PbO,. The relation is clearer 
if we write the formula PbePbO.. When this is treated with 
nitric acid, the two lead atoms which form the positive ions, 
Pbt*, of the salt, are easily exchanged for the hydrogen of the 
acid, but the plumbie acid, H4PbO,, which results, is unstable 
and decomposes at once to water and plumbic anhydride or 
lead dioxide, PbOs, just as carbonic acid, H2CO3, decomposes 
to carbon dioxide and water. 

Red lead is used as a pigment, as an oxidizing agent in glass 
manufacture and with linseed oil as a lute in plumbing. 

As litharge is oxidized to lead plumbate, Pb2PbOu,, a mixture of 
lime, CaO, and litharge may be readily oxidized at low redness 
to calcium orthoplumbate, CagPbO,. Dilute acids decompose 
this with the separation of lead peroxide. 

Storage Batteries. A storage battery which has been charged 
contains two kinds of lead, one of which has been more or less 
completely changed to lead dioxide, while the other consists 
of spongy, metallic lead. During the discharge the electrons 
escape to the connecting wire from the metallic lead, Pb, chang- 
ing it to lead ions, Pb**, which combine with sulfate ions of the 
solution, forming insoluble lead sulfate, PbSO,: 


Nid is ote CF wame el wen OV ES = oie 


The electrons pass through the plate and connections to the 
other plate. At the other plate the two electrons combine with 
the lead of the lead dioxide, reducing it from the quadrivalent 
to the bivalent form, resulting in the formation of lead sulfate 
and the liberation of a sulfate ion, SO,-~, from the sulfuric 
acid : 


2- + PbO, + 2 H.SO, = PbSO. + 2 H2O + SO.-— 


STORAGE BATTERIES 517 


or Boe teOnnOse bee] PbeszOne 4+ O-7 
O-~ + H.SO, = H.O + SO 
and Pb*+O-~- + H.SO, = Pbt*SO,= + H,O 


During the discharge there is a difference of potential of 
about two volts between the two plates, and the passage of the 
current develops a very considerable amount of electrical 
energy. In charging the battery the reverse operations take 
place. At the cathode, which is connected with the negative 
pole of the dynamo, electrons combine with lead ions, Pbt*, 
reducing them to metallic lead, Pb, and leaving the sulfate ions, 
SO,-~, of the lead sulfate free to pass into solution. 


Pb*t+80.-7> + 27> = Pb + S0,-- 


At the anode the bivalent lead ions, Pb*t, lose two electrons, 
giving tetravalent lead ions, Pb+*++t, which combine momentarily 
with another sulfate ion, SO.-~, to form lead tetrasulfate, 
Pb(SO,)2. The lead tetrasulfate is at once hydrolyzed to lead 
peroxide, PbOs, and sulfuric acid. 


Po Om sO, Pb 0. 
Pb(SO.)2 + 2H2O = PbO, + 2H.SO, 


It will be noticed that in both processes sulfate ions, SOs, 
are discharged at one plate and enter into combination with 
the lead, while at the other plate sulfate ions pass into solution. 
In charging the battery the discharged sulfate ions combine 
with the lead of lead sulfate, PbSO., forming the tetrasulfate, 
Pb(SO,)2. In discharging the discharged sulfate ions combine 
with the metallic lead, forming lead sulfate. At the same time 
sulfate ions must, of course, migrate through the solution in one 
direction in charging, in the opposite direction in discharging. 

The charging is, of course, accompanied by an_ absorp- 
tion of energy. Practically, a very high efficiency can be se- 
cured, the energy obtained during the discharge approaching 
closely to that absorbed in charging. Since the sulfate radicals 
are mostly in the form of sulfuric acid in the charged battery 


518 A TEXTBOOK OF CHEMISTRY 


and in the form of lead sulfate in the battery after discharge, 
the amount of sulfuric acid in the electrolyte, which can be easily 
determined by a hydrometer, furnishes a pretty close indication 
of the condition of the cell. 

Lead dioxide when warmed with hydrochloric acid gives at 
first lead tetrachloride, PbCli, but this is unstable and decom- 
poses into lead chloride, PbCl:, and chlorine, Cl, exactly as 
manganese tetrachloride does (p. 101). It is noteworthy that 
this is entirely different from the conduct of barium peroxide, 
BaOz, or sodium peroxide, Na2Ox, either of which gives hydrogen 
peroxide, H.O2, with hydrochloric acid. (What does this dif- 
ference in conduct indicate as to the structure of these three 
oxides ”) | 

Lead Sulfide, PbS, is found in nature as the mineral galena 
and is formed as a black precipitate by the action of hydrogen 
sulfide on a solution of a lead salt in a dilute acid. It is not 
precipitated in the presence of much hydrochloric acid, or of 
much of any other strong acid. It dissolves easily in nitric 
acid. 

Lead Chloride, PbCl2, forms as a white, crystalline precipi- 
tate on adding hydrochloric acid to a solution of almost any 
soluble salt of lead. It dissolves in 125 parts of water at 18° 
and in 30 parts of boiling water. It is less soluble in dilute 
hydrochloric acid than in pure water, but is more soluble in 
concentrated acid, doubtless because of the formation of a 
complex compound with the acid, such as chloroplumbous acid, 
HePbCl,. 

Lead Tetrachloride, PbCl, By dissolving lead peroxide, 
PbOs, in cold, concentrated hydrochloric acid a solution con- 
taining lead tetrachloride, or more likely chloroplumbic acid, 
H2PbCls, is obtained. On the addition of ammonium chloride, 
ammonium chloroplumbate, (NH,4)2PbCle, separates. If this 
salt is dissolved in cold concentrated sulfuric acid, hydrochloric 
acid escapes and lead tetrachloride separates below the acid as a 
heavy yellow liquid. It is quite unstable, decomposing readily 
into lead chloride, PbCl2, and chlorine. It is hydrolyzed by 


LEAD SALTS 519 


water to lead dioxide and hydrochloric acid. With a little 
hydrochloric acid it gives a yellow compound, chloroplumbic 
acid, HePbCl.e. A considerable number of salts of this acid are 
known. The ammonium salt is mentioned above. These salts 
are usually called double salts of lead tetrachloride with other 
chlorides, and are frequently written in the form PbCl,.2 KCl 
instead of K2PbClg. 

Lead Sulfate, PbSO., may be eras ah by the precipitation 
of any soluble salt of lead with dilute sulfuric acid. It may 
also be obtained by roasting lead sulfide at a moderate tem- 
perature. The compound prepared in the latter manner has 
been used in America to a limited extent as a pigment in place 
of the ordinary white lead described below. It is a difficultly 
soluble salt, but is distinctly more soluble than barium sulfate. 
It dissolves easily in a solution of ammonium acetate. 

Lead Nitrate, Pb(NOs)2, is a very easily soluble salt which 
may be prepared by dissolving either lead or litharge in dilute 
nitric acid. Either lead nitrate or lead acetate may be used 
for the preparation of the insoluble salts of lead, especially of 
the chromate, PbCrQOx.. 

Lead Acetate or Sugar of Lead, Pb(C2H302)2.3 H2O, is an 
easily soluble salt prepared by dissolving litharge in acetic 
acid. Sugar of lead has sometimes been used in hair dyes, but 
its use in this way is considered dangerous and liable to cause 
paralysis. 

Basic Lead Acetates, Pb(C2H;02)OH and Pb(C2H3O2)e. 
2 Pb(OH)s, are formed by dissolving litharge, PbO, in a solu- 
tion of lead acetate. Such a solution is used to clarify dark- 
colored sugar solutions to prepare them for determinations with 
the polarimeter. 

Lead Carbonate, PbCO;. The normal salt may be precipi- 
tated by adding sodium carbonate to a solution of lead acetate. 
It is so difficultly soluble that it can also be precipitated by 
passing carbon dioxide through a solution of lead acetate, and 
‘lead acetate gives a turbid solution with ordinary distilled water 
because of its formation. 


520 A TEXTBOOK OF CHEMISTRY 


Basic Lead Carbonate, or White Lead, Pb3(CO3)2(OH).2 or 
2 PbCO;3.Pb(OH)s, has been manufactured for many centuries 
for use as a pigment. ‘The principal process used, known as the 
“Dutch process,” has been scarcely changed in principle for a 
very long time. , 

Plates, or “ buckles,” of lead about one eighth of an inch 
thick and five and one half inches in diameter are cast in the 
form shown in Fig. 104. These plates are packed in pots, Fig. 
105, having about 250 cc. of dilute acetic acid or vinegar in the 





Fig. 105 


bottom. These pots are then packed in layers with alternate 
layers of spent tanbark until a large room is filled. The room 
is left to itself for about three months. ‘The combined action 
of the vapors of acetic acid and air on the lead plates causes 
them to corrode superficially, and the basic acetate formed is 
changed to carbonate by the carbon dioxide which comes from 
the fermentation of the tanbark. As the carbonate is formed 
some acetic acid is continuously liberated, and this, with the 
air, carries on the corrosion till the ‘‘ buckles”? are almost 
completely changed to basic carbonate. At the end of three 
months the pots are emptied, the white lead is finely ground 
and separated from particles of lead and coarse particles of 
material by bolting and lixiviating with water. The fine powder 
is then dried and intimately incorporated with linseed oil for 


WHITE LEAD ° 521 


the market. Another process is to mix the moist powder 
directly with linseed oil. Owing to the relation between the 
surface tension of white lead toward water and that toward 
linseed oil, the latter is able to displace the water and joins with 
the white lead to form a paint which is practically identical in 
composition and properties with the pigment prepared from the 
dry powder. Owing to the poisonous character of lead com- 
pounds, workmen in white lead factories and painters often 
suffer from a painful and sometimes fatal disease, called lead 
colic. Stringent laws have been passed by some states for the 
protection of the workmen from inhaling the dust and from 
poisoning in other ways. Some manufacturers spend large sums 
of money to protect their workmen from the danger of poisoning. 

White lead depends for its value on the fact that it is an 
amorphous, very fine, opaque white powder, the opacity being 
much greater than that of barium sulfate, which is sometimes 
used as an adulterant or substitute. The fact that water will 
not wet it when it is in contact with linseed oil is also a factor 
of prime importance. White lead is blackened by hydrogen 
sulfide and for that reason is less suitable than zine white and 
lithopone for use in chemical laboratories or in localities where 
it is subjected to the action of sewer gas. 


CHAPTER XXX 


VANADIUM AND CHROMIUM GROUPS 


Group V. Vanadium, columbium (or niobium) and tantalum 
are the elements of Group V, which alternate with phosphorus, 
arsenic, antimony and bismuth. They are usually classed as 
rare elements, but vanadium and tantalum have acquired some 
industrial importance. Each forms a pentoxide, corresponding 
to POs, and salts of acids corresponding to the metaphosphates, 
MPO, pyrophosphates, M,P20,7, and orthophosphates, M3POsx. 
Salts of the form MgV2Oy, are also known for each. 

Vanadium, V, 50.96, is very widely diffused in nature, being 
found in small amounts in almost all clays and massive rocks. 
The most important mineral is vanadinite, Pb;(VO.)3Cl, which 
corresponds in composition to apatite, Cas(PO.)sF. 

Vanadium is a silvery white metal having a specific gravity 
of 5.5 and melting at about 1720°. It is used as an addition to 
steel, increasing its hardness, malleability and tensile strength. 
Ferrovanadium is an alloy with iron, which is much more easily 
prepared than the pure metal. It is the commercial form used 
for addition to steel. 

* Vanadium forms compounds in which it is bivalent, tri- 
valent, quadrivalent and quinquivalent. The following may 
be mentioned: vanadous chloride, VCl2, vanadous sulfate, 
VSO,.7 H,O, vanadous sulfide, VS, vanadie chloride, VCl3, 
vanadic sulfide, V2S3, vanadie sulfate, Vo(SOx4)3, vanadium alum, 
KV(SOu4)o.12 H2O, vanadium tetrachloride, VCli, vanadium 
pentachloride, VCl;, vanadium oxychloride, VOCI]3, sodium 
orthovanadate, Na3VO..12H:O, sodium  pyrovanadate, 
NaeV20;7.18H.O, sodium metavanadate, NaVO3.2H.O. More 
complex vanadates and other complex compounds of a great 


variety of forms have been prepared. 
522 


COLUMBIUM. TANTALUM O24 


* Columbium, Cb, 93.1 (or Niobium, Nb). In 1801 Hatchett 
discovered a new element in a mineral from Haddam, Con- 
necticut. He called the mineral columbite (from Columbia, 
the poetical name for America), and the element columbium. 
It is probable that the compounds which he prepared contained 
both columbium and tantalum, and the two elements were first 
clearly séparated and characterized by H. Rose in 1844. Rose 
either overlooked or ignored the discovery of Hatchett and 
called the two metals tantalum and niobium. 

Columbite, the mineral in which columbium was discovered, 
is a ferrous metacolumbate, Fe(CbO3)2, containing ferrous 
metatantalate, Fe(TaQOs3)2, in isomorphous mixture. The 
formula is more properly written, Fe((Cb,Ta)O3) or, in 
the oxide form, which is most frequently used by mineralogists, 
FeO((Cbe2,Taz)Os). 

Elementary columbium is still more metallic in its properties 
than vanadium. It has a specific gravity of 7.4 and melts at 
about 2200°. The following are typical compounds: colum- 
bium trichloride, CbCl;, columbium pentafluoride, CbF;, colum- 
bium oxyfluoride, CbOF;, columbium pentoxide, Cb2O;, mag- 
nesium orthocolumbate, Mg3(CbO.)2, calcium pyrocolumbate, 
CasCb2O7, potassium hexacolumbate, 4 K20.3 Cb.0;.16 H2O or 
KsCb209.16 H20. Many complex columbates ana other com- 
plex compounds are known. 

* Tantalum, Ta, 181.5. In 1902 Ekeberg examined a mineral 
from Finland which closely resembles columbite and gave to 
it the name tantalite. He called the element which it contains 
tantalum, but the compounds which he prepared were doubtless 
mixtures containing both columbium and tantalum. The 
mineral which he studied always contains both elements, and 
it is most properly called columbite when the columbium is in 
excess and tantalite when there is more of the tantalum. As 
stated above, H. Rose first distinguished sharply between the 
two elements and cleared up the confusion of the earlier workers. 

Tantalum is a bright metal somewhat resembling platinum 
in appearance. It is ductile and can be drawn into fine wire. 


924 A TEXTBOOK OF CHEMISTRY 


As the melting point is 2850° it is suitable for the filaments of 
electric lights, and was used for a short time in that way, but 
was quickly displaced by tungsten. The specific gravity of the 
metal is 16.5. 

Among the compounds of tantalum are the potassium fluo- 
tantalate, 2KF'.TaF; or KeTaF;, by means of which tantalum 
can be best separated from columbium (E. F. Smith), tantalum 
pentachloride, TaCl;, tantalum pentoxide, Ta.O;, and a series 
of tantalates. 

Group VI. Chromium, molybdenum, tungsten and uranium 
alternate with sulfur, selenium and tellurium of Group VI, 
exactly as the three elements last considered alternate with the 
elements of the phosphorus family. 

The elements form compounds in which they appear bivalent, 
trivalent and sexivalent, especially, but also some in which 
they are quadrivalent. They resemble the sulfur family in 
such compounds as potassium chromate, K»2CrQ,., potassium 
dichromate, K»Cr.0;, and ammonium molybdate, (NH4)2MoO.. 

Chromium, Cr, 52.01, is found chiefly as ferrous chromite, 
FeO.Cr203 or Fe(CrO2)2, which is known as chromite or chrome 
iron ore, a black mineral isomorphous with magnetite, Fe;O, or 
(FeO.Fe.03). Chromium is also found as lead chromate, 
PbCrOu., and was first discovered in that mineral by Vauquelin 
in 1797. The name was given because of the colored compounds 
which it forms. 

Metallurgy, Uses. Chromium is now prepared by Gold- 
schmidt’s thermite process (p. 497) by igniting a mixture of 
chromic oxide, CreO3, and aluminium. 

Metallic chromium is a white, crystalline, extremely hard 
metal — as hard as corundum. It has a specific gravity of 
about 7.0. It melts at 1520° and boils at 2200°. It is used as 
an addition to steel, making it extremely hard and resistant to 
the penetration of projectiles when used for armor plate. 
Such steel requires very careful heat treatment to bring out its 
best properties. Alloys with nickel or cobalt resist the action 
of acids and are proving useful for special purposes. 


CHROMIUM 525 


Chromous Chloride, CrCly..4H,O. A light blue solution 
containing chromous chloride is easily prepared by the action 
of zine and hydrochloric acid on a solution of chromic chloride, 
CrCl;, or of chromiv anhydride, CrO3. Air must be carefully 
excluded, as air or oxygen, in the presence of hydrochloric acid, 
changes the compound back to green chromic chloride. The 
solid chromous chloride is white. 

Chromic Oxide, Cr2O3;, is a green powder formed by igniting 
the hydroxide. A more or less pure chromic oxide or hydroxide 
is prepared in a variety of ways and used as a pigment under 
the name of chrome green or Guignet’s green. The latter con- 
tains a little boric acid. A green pigment which is made by 
mixing chrome yellow, PbCrQu,, and Prussian blue, Fea(FeCgNe)s, 
is also often sold and used under the name of “‘ chrome green.” 

Chromic Hydroxide, Cr203.4 H.O. From analogy with the 
formula of the chloride, the formula Cr(OH)3 is often given for 
the hydroxide. The compound formed by precipitating a 
chromic salt with ammonium hydroxide has the composition 
Cr203.4 HO, however, when it is dried in a vacuum. A com- 
pound, or mixture, having the composition Cr(OH)s; has been 
obtained in some cases, but it is not the usual composition of the 
hydroxide. Chromic hydroxide dissolves to some extent in 
alkalies, but does not show as marked acidic properties of 
aluminium hydroxide. A number of chromites have been pre- 
pared, however, and ferrous chromite, Fe(CrO2)s, has been men- 
tioned above as the most important ore of chromium. 

Chromic Chloride, CrCl:, was formerly prepared by heating 
a mixture of chromic oxide, Cr2O3, and carbon in a current of 
chlorine. It sublimes in beautiful reddish pink or violet leaflets. 
These are almost insoluble in water, but dissolve easily to a green 
solution in water containing a trace of chromous chloride, CrCly. 

Hydrates of Chromic Chloride. A _ solution containing 
chromic chloride is easily prepared by the reduction of chromic 
anhydride or a solution of a chromate: 


2 CrO;3 + 6 HCl + 3C.H;OH = 2? CrCls + 3 C.H,O + 6 H2O 
Alcohol Aldehyde 


526 A TEXTBOOK OF CHEMISTRY 


Other reducing agents, such as sulfurous acid, may be used, 
but alcohol has the advantage that it leaves no other compound 
in the solution. 

By various methods it is possible to prepare two isomeric 
hydrates of chromic chloride, both of which have, in the erystal- 
line form, a composition corresponding to the formula 
CrCl3.6 H,O. One of these hydrates dissolves in water to a 
green solution, while the other gives a violet solution. Each 
can be changed more or less completely into the other, and 
solutions of chromic chloride usually contain the two compounds 
in equilibrium with each other. ‘The green solution has a lower 
electrical conductivity than the violet. If silver nitrate is added 
to an ice-cold solution of the green hexahydrate, only one third 
of the chlorine is precipitated at first. 

The most satisfactory explanation of these facts is given by 
the theory of Werner. He supposes that in each compound six 
atoms or groups are directly combined with or arranged about 
the chromium atom to form a characteristic group. In the 
violet chloride he supposes that these six groups are all mole- 
cules of water and he calls this hydrate hexaaquochromic chloride 
and writes the formula (Cr(OH:2).)Cl3. In the green chloride, 
on the other hand, he considers that four molecules of water and 
two atoms of chlorine form the six groups directly combined 
with the chromium, while one atom of chlorine and two mole- 
cules of water are less directly connected. He writes the formula 

Cl, 

a Cl + 2 H.O. 
Three chlorine atoms of the violet chloride are readily 
ionized, giving electrical conductivity and easy precipita- 
tion as silver chloride. But only one of the chlorine atoms 
of the green chloride ionizes to give electrical conductivity, 
and only this one atom is directly precipitated by silver 
nitrate. 

Werner and others have shown that a large number of com- 
plex inorganic compounds, containing water, ammonia and 
other groups, exhibit isomerism similar to that of these hydrates. 


of the green chloride, accordingly, [cr 


CHROMATES 527 


(See Werner, Neuere Anschauung auf dem Gebiet der anorgan- 
ischen Chemie.) 

Potassium Chromium Sulfate or ‘Chrome Alum,” 
KCr(SOx4)2.12 HO. If potassium dichromate, K2Cr2O;, is 
warmed with alcohol and dilute sulfuric acid, the chromium is 
reduced to the form of chromic sulfate, Cre(SOx4)3, and on evapo- 
ration of the solution and cooling, the chromic sulfate and potas- 
sium sulfate will combine to form chrome alum, a dark violet 
salt, which is isomorphous with ordinary alum. 

Potassium Chromate, K.CrO,;. When chrome iron ore, 
FeCr.0u,, is roasted with potassium carbonate, the chromium is 
oxidized and forms potassium chromate. This is in accordance 
with the practically universal rule that metals assume a higher 
state of oxidation in the presence of a base than in the presence 
of an acid: 


4 FeCr.O, + 8 K.COs3 ste 7 O. ty Fe.O; + 8 KeCrO, + 8 CO, 


The potassium chromate is easily soluble and can be easily 
separated from the insoluble ferric oxide. It is an easily soluble 
yellow salt, which gives a lemon-yellow solution. 

Potassium Dichromate, or Pyrochromate, K,Cr.O7, is formed 
on adding an acid, even a weak acid, to a solution of potassium 
chromate. It crystallizes in orange-red crystals, which are 
much less soluble than the crystals of the chromate. Potassium 
dichromate is the practical source of almost all other chromium 
compounds and exceeds all of the others in commercial impor- 
tance. It is used as an oxidizing agent, as a mordant in dyeing, 
and for the preparation of leather in chrome tanning, a process 
rapidly increasing in importance. 

Potassium dichromate corresponds to potassium pyrosulfate, 
K.S,0;. The chromium in it is in the same state of oxidation 
as in the chromate. 

Lead Chromate, or Chrome Yellow, PbCrQ,, is an insoluble salt 
obtained by precipitating a solution of lead acetate with potassium 
dichromate. It is a brilliant yellow compound which is used as 
a pigment and as a constituent of ‘‘ chrome green ”’ (p. 525). 


528 A TEXTBOOK OF CHEMISTRY 


* Barium Chromate, BaCrQs,, is easily prepared by precipita- 
tion. It is insoluble in water or acetic acid but dissolves in hy- 
drochloric or nitric acid. 

Chromium Trioxide or Chromic Anhydride, CrO;, separates 
in the form of dark red needles on adding an excess of concen- 
trated sulfuric acid to a saturated solution of potassium 
dichromate. When mixed with sulfuric or other acids, it is a 
powerful oxidizing agent and is much used for that purpose 
in the laboratory. To illustrate; such a mixture with sulfuric 
acid will oxidize the carbon or graphite of cast iron or steel to 
carbon dioxide, and this method is often used for the determi- 
nation of these elements. 

Chromyl Chloride, CrO2Cle, is prepared by distilling a mixture 
of potassium dichromate, salt and sulfuric acid. It is a volatile, 
fuming, dark red liquid, which is a very powerful oxidizing 
agent. Many organic compounds react with it with explosive 


| H—O O 

violence. It may be considered as chromic acid, yak . 
H—O O 

in which the two hydroxyl groups have been replaced by chlorine, 


oe f? 
Dat . It is readily hydrolyzed by water to hydrochloric 
a 2 O OH 
Yor’ 
OF 


\ 


and dichromic acids. The later may be written ips : 


“ok 
Oar 

Molybdenum, Mo, 96, occurs chiefly as molybdenite, Moss, 
a black mineral closely resembling graphite in appearance but 
having more than twice the specific gravity of that mineral. 
The metal is silver-white, after melting, and has a crystalline 
fracture. It melts at about 2500°. 

* Molybdium Trioxide, or Molybdic Anhydride, MoOs, is a 
yellowish white or white powder, which dissolves easily in alka- 
lies, forming molybdates, such as sodium molybdate, NasMoO,, 


MOLYBDENUM 529 


It is nearly insoluble in acids, but if a solution of ammonium 
molybdate, (NH4)2MoOs,, is poured into dilute nitric acid (not 
the reverse), the molybdic acid or anhydride which is formed 
remains in solution either as a colloid or in a supersaturated solu- 
tion. Such a solution, known as “‘ molybdic solution,” is used 
to precipitate solutions of orthophosphoric acid, H;PO,. On 
adding an excess of the molybdic solution to a solution contain- 
ing a soluble orthophosphate, or orthophosphoric acid, a very 
difficultly soluble, yellow precipitate separates. This has the 
composition (NH,)3PO4.12 MoO3.nH2O. As the molybdic an- 
hydride will neutralize alkalies, the amount of the molybde- 
num and so, indirectly, that of the phosphorus may be de- 
termined by titration with a standard solution of potassium 
hydroxide : 


bie) 


MoO; + 2 KOH = K;Mo0O, + H,O 


If the precipitate is dissolved in ammonia and an excess of dilute 
sulfuric acid added, the molybdenum may be reduced to molybdiec 
sulfate, Moe(SO4)3, by means of zine. The solution obtained in 
this manner may be oxidized quantitatively to a solution con- 
taining molybdie anhydride by potassium permanganate. This, 
again, gives an indirect determination of the amount of phos- 
phorus. The solution of molybdic sulfate is very easily oxi- 
dized by the air and the success of the operation depends on rapid 
work. 

* Compounds of Molybdenum. Molybdenum forms a series 
of oxides, the most important or best characterized being, Mo2Qs, 
Mo;O12, Mo20;, Mo303, MoO2z and MoQO3. The sulfides are 
Mo,8s3, Mos:,, MoS; and MoS.. 

* Molybdic anhydride, MoOs, forms a bewildering variety of 
complex compounds of which the ammonium phosphomolyb- 
date used in analysis and referred to above is an isolated exam- 
ple. Among the related compounds are many ammonium 
molybdates, such as_ triammonium  dodekamolybdate 
12 MoO3.3 NH3.12 H,O, and phosphomolybdie acids, as phos- 
phoduodecimolybdie acid, 24 MoO3.P205;.4 H20O. 


530 A TEXTBOOK OF CHEMISTRY 


A number of different ammonium phosphomolybdates have 
also been prepared, and it is only by careful attention to the 
proper conditions of temperature, acidity and concentration of 
the reacting solutions that a compound having the composition 
given above can be obtained. It is much easier to be sure that 
all of the phosphoric acid is precipitated from the solution than 
to secure the correct ratio between molybdic anhydride and phos- 
phoric acid. For this reason it is a common analytical practice 
to dissolve the precipitate in ammonia and precipitate the phos- 
phorie acid as magnesium ammonium phosphate, MgNH,.PO,. 

Tungsten, W, 184.0. Although tungsten is usually classed 
among the less common elements, its compounds have been 
known since the middle of the eighteenth century. As a constit- 
uent of the old ‘‘ Damascus blade ”’ its valuable effect on steel 
was used long before its presence was recognized, and some of its 
compounds have long been used for fireproofing fabrics. The 
recent use for the filaments of electric lights has now made the 
name of the metal a household word. 

Tungsten is found chiefly in the form of wolframite, a ferrous 
manganese tungstate, (FeMn)WO,, the iron and manganese 
replacing each other as isomorphous constituents. Metallic 
tungsten can be obtained by the reduction of tungstic anhydride, 
WOs, with carbon or hydrogen at a high temperature, or by 
Goldschmidt’s thermite process. 

Tungsten is a heavy, steel-gray, very hard metal. A good deal 
of difficulty was experienced in learning how to draw the metal 
into wire suitable for incandescent electric lights. The specific 
gravity is 18.64. Tungsten melts at 3000°, the highest melting 
point of any element except carbon. Its use for electric lights 
depends, of course, on this property. The temperature for rapid 
volatilization is probably higher than that of carbon. 

A very important application of tungsten is in the manufacture 
of “ high-speed tool steel.”’ Ordinary steel cannot be used for 
any rapid lathe work because it would become so hot as to lose 
its temper. Some kinds of tungsten steel are very hard and will 
also retain their hardness even when almost red-hot. The 


TUNGSTEN. URANIUM Tail 


introduction of such tools has almost revolutionized shop 
practice in America. 

* Compounds of Tungsten. The element forms four chlo- 
rides: tungsten dichloride, WCl2, tungsten tetrachloride, WCl,, 
tungsten pentachloride, WCI;, and tungsten hexachloride, WClg. 
The last two are volatile, while the dichloride and tetrachloride 
are not. The two most important oxides of tungsten are the di- 
oxide, WOs, and the trioxide, or tungstic anhydride, WO;3. Very 
many complex tungstates have been prepared. In addition to the 
normal sodium tungstate, Na2zWO,, no less than thirteen complex 
sodium tungstates containing a smaller proportion of tungsten 
have been described. The compound 5 Na2,O.12 WO; may be 
mentioned as an illustration of these compounds. Sodium 
tungstate has been used as a mordant in dyeing and for the 
fireproofing of fabrics. 

By the reduction of acid sodium tungstates by heating them 
with tin or hydrogen a series of yellow, blue, violet and purple 
compounds called ‘‘ bronzes’ has been prepared. These con- 
tain less oxygen than true tungstates should. Thus the composi- 
tion Na,W.O, is given for the violet bronze, while the corre- 
sponding tungstate would be NazW.4013 = Na2O.4 WOs. 

* Phosphotungstic Acid, H3PO.4.12 WO3.18 H.O. This com- 
pound can be prepared by treating silver tungstate, AgoWOu,, 
mixed with the calculated amount of phosphoric acid, with 
hydrochloric acid. It crystallizes in rhombic crystals which are 
solubie in water. The solution gives characteristic precipitates 
with alkaloids and with proteins and is used as a reagent for 
these purposes. Several other complex phosphotungstic acids 
have been prepared. 

Uranium, U, 238.17. Uranium is found chiefly in the form of 
uraninite, or pitchblende, or U3;Os. Several other minerals con- 
tain uranium and all of these are now of interest because of the 
connection with radium (p. 471). Cleveite, a mineral from 
Norway, resembles pitchblende, but contains also yttrium and 
other rare elements. It is the mineral in which Ramsay first 
discovered helium on the earth (p. 237). Carnotite is a uranate 


032 A TEXTBOOK OF CHEMISTRY 


and vanadate of potassium K,0.2 UO3.V20;.3 H2,O. Samars- 
kite is a complex tantalo-columbate of uranium, yttrium, iron 
and other metals. 

Uranium is a white metal. When free from carbon it is not so 
hard as steel. Its specific gravity is 18.68. It melts at a higher 
temperature than platinum. 

The oxides of uranium are uranium dioxide, UOs, a green 
oxide, U30g, of the same composition as uraninite and uranium 
trioxide or uranic anhydride, UO3;. The chlorides are uranium 
trichloride, UCls, uranium tetrachloride, UCls, uranium penta- 
chloride, UCl;, and uranium hexachloride, UCIlg. 

Comparatively few salts of uranium are known in which metal- 
lic uranium replaces the hydrogen of an acid directly. In these 
few it is quadrivalent. One of the simplest is uranium sulfate, 
U(SO,4)o.2 H2O. The more common salts of uranium contain 
the bivalent group uranyl, UO». Thus uranyl nitrate is 
UO2(NOs)2.8 H2O and uranyl acetate is UOe(C2H302)2.2 H2O. 

The compounds in which uranium acts as an acid forming ele- 
ment are mostly diuranates, corresponding to the dichromates. 
Potassium diuranate, K2U:2O;, is an orange-yellow, almost in- 
soluble powder. 

As a radioactive element uranium has a “ half-life’ of about 
6,000,000,000 years. 


CHAPTER XXXI 


MANGANESE 


Group VII. While there are three elements (V, Cb, Ta) al- 
ternating with phosphorus, arsenic, antimony and bismuth in 
Group V, and four elements (Cr, Mo, W, U) alternating with 
sulfur, selenium and tellurium in Group VI, manganese is the 
only element alternating with chlorine, bromine and iodine in 
Group VII. Not only is the halogen of atomic weight about 
214 missing, but elements resembling manganese with atomic 
weights approximately 98, 187 and 242 have never been found. 
The radioactivity of uranium and the ephemeral life of radon 
(p. 474) which are found in this region of the table have recently 
given us a hint as to a possible reason for these gaps in the sys- 
tem. It seems likely that the structure of the atom which would 
give elements of these atomic weights is unstable, and that either 
these elements cannot exist at all or they are to be looked for 
among the radioactive elements of brief life. 

Manganese has properties such as its position in the table 
leads us to expect. As a metal it closely resembles iron, cobalt 
and nickel, all four elements forming a transition from the hard, 
difficultly fusible chromium to copper with its much lower melt- 
ing point (1083°) and its great malleability and ductility. In its 
nonmetallic properties, on the other hand, it forms acidic oxides 
and acids, and the highest of these, permanganic acid, HMnOQOu,, 
corresponds to perchloric acid. 

Manganese, Mn, 54.93. Occurrence, Properties. Manganese 
is found chiefly as the dioxide in the mineral pyrolusite, MnOx. 
It is also found in small amounts in most minerals and rocks, in 
practically all iron ores and in some natural waters. Except 
for scientific purposes the element is not prepared in the pure 


state. Pure manganese is a very hard, reddish gray metal with 
533 


534 A TEXTBOOK OF CHEMISTRY 


a specific gravity of 7.2. It melts at 1260°. It dissolves easily 
in acids, even more easily than iron, forming manganous SN 
in which the element is bivalent. 

Alloys of manganese with iron are easily manufactured, com- 
mercially, by reducing a mixture of the ores of the two metals 
in a blast furnace (p. 541). Those containing 10-15 per cent 
of manganese are white, with a brilliant metallic luster, and re- 
tain the carbon in the combined form. Because of the appear- 
ance of the surface, this alloy is called spiegeleisen (mirror-iron). 
It is used in the manufacture of steel (p. 548). An alloy con- 
taining 70-90 per cent of manganese with iron is called ferro- 
manganese and is used as an addition to cast iron. Manganese 
bronze is an alloy of manganese and copper containing 30 per 
cent of manganese. It is hard and has a high tensile strength. 

Compounds of Manganese. In compounds in which it acts as 
a metal, manganese is almost exclusively bivalent, as in manga- 
nous chloride, MnCl, and manganous sulfate, MnSQ,. There 
are a few unstable compounds, such as manganic chloride, MnCl, 
in which it is trivalent ; and compounds in which it is a quadri- 
valent, basic element have been prepared only in the form of 
double salts such as 2 KC].MnClu, or K2MnCl,s. The basic 
oxides are manganous oxide, MnO, and manganese sesquioxide, 
Mn:,Os. 

As an acid-forming element manganese is quadrivalent in 
manganese oe Mn0O,, and the manganites, such as calcium 


manganite, Ca "SMn=0, It seems to be sexivalent in the 
O 
Kk a ype 
manganates, such as potassium manganate, Mné 
K-- Of sais: 


or K,.MnQ,, and septivalent in the bapancere such as 


Ws 
potassium permanganate, K— heya 
Sy 

Manganous Manganic Oxide, Mn;QO,, is of a mixed type, some- 
what similar to that of red lead, but the similarity in formulas 


MANGANESE 535 


is probably superficial. Red lead is lead plumbate, Pb:PbO,, 


fares 

or pone while the oxide, Mn;O,, is most likely derived 
nga 

from a hypothetical acid HMnO.. _ If this is true, the structure 


Oo aein © 
is Mn(MnO,), or Mnx 
\o—Mn=O 


* Manganous Hydroxide, Mn(OH)., is a white precipitate 
which quickly turns brown from oxidation on exposure to the air, 
especially if an alkali is present. 

* Manganous Chloride, MnCl:.4 H,0, is a light pink, easily 
soluble salt. 

* Manganous Sulfate, MnSOu,, forms hydrates with 1, 2, 3, 4, 
5 or 7 molecules of water. The last, which corresponds in com- 
position to white vitriol, ZnSO4.7 H.O, and green vitriol or 
eopperas, FeSO..7 H.O, can only be obtained by crystalliza- 
tion at temperatures below 6°. 

* Manganous Sulfide, MnS. Ammonium sulfide precipitates 
from alkaline solutions of manganous salts a flesh-colored precip- 


SH 
itate having the composition MnS.H,0 or Ming . Itis easily 
OH 
soluble in acids, even in acetic acid. The sulfide, MnS, is an 
olive green powder which is formed by the action of hydrogen 
sulfide on any of the oxides. 

Manganese Dioxide, MnOnz, or Black Oxide of Manganese is 
found as the mineral pyrolusite in sufficient quantities to make it 
a very valuable commercial product. The compound has played 
a commanding roéle in the development of chemical knowledge 
and also in a variety of industrial processes. By means of it 
Scheele discovered chlorine, and it was early used in the prepara- 
tion of oxygen. It is still constantly used in laboratories to 
catalyze the decomposition of potassium chlorate (p. 21). 

During the latter half of the nineteenth century manganese 


536 A TEXTBOOK OF CHEMISTRY 


dioxide was used in large quantities for the preparation of chlo- 
rine as an adjunct to the Leblanc soda process (p. 411). When 
the supply of manganese ore declined and the material grew ex- 
pensive, the Weldon process (p. 103) was introduced to conserve 
the manganese. At the present time, when electrolytic sodium 
hydroxide and chlorine are promising to complete the extinction 
of Leblanc soda and compete strongly with ammonia-soda, man- 
ganese is finding increasing use in the metallurgy of iron and 
steel, and iron manufacturers are searching eagerly for new sup- 
plies of the ore. 

It is worthy of notice that the Weldon process (p. 103) depends 
on the tendency of the metallic elements to assume a higher 
valence toward oxygen, especially in an alkaline solution, than 
their usual valence toward chlorine and other acid radicals in 
an acid solution. The sign of the valence does not, however, 
change. Manganese seems to be positive in all of its com- 
pounds. 

Manganates. When manganese dioxide, or indeed almost 
any compound of manganese, is fused with potassium or sodium 
carbonate and some oxidizing agent, as potassium nitrate or 
potassium chlorate, potassium manganate, K2MnQ,, or sodium 
manganate, NasMnQOu,, is formed. This has a green color and 
dissolves in water to a dark green solution. The solution is 
sometimes called “ chameleon solution’ because of the color 
changes which it undergoes so easily as the manganate changes 
to red permanganate or turns brown from the separation of man- 
ganese dioxide. The manganates are stable only in an alkaline 
solution, and free manganic acid, HeMnQu,, does not, apparently, 
exist, even in solution. The addition of an acid, even of carbonic 
acid, to the solution causes an autoxidation and reduction some- 
what similar to that of hypochlorous acid (p. 127) or to that of 
potassium chlorate to potassium perchlorate and potassium chlo- 
ride (p. 128). 


H.MnO, -f 2 H.MnO, ap HMnO, + MnO, + Z H.O 


Manganic Permanganic 


Acid Acid 


PERMANGANATES Dod 


(In developing this equation, notice, (1) that MnO.+ H,0 is 
equivalent to H,MnQOs, hence one molecule of manganic acid 
gives one atom of available oxygen, and (2) that one atom of 
oxygen will oxidize two molecules of manganic acid to perman- 
ganic acid by the removal of two atoms of hydrogen.) 

Permanganates. Manganic acid is a very weak acid and its 
salts are strongly hydrolyzed in the absence of an alkali: 


K.MnO, ot 2 HOH = H.MnO, + 2 KOH 


As has just been stated, the manganic acid formed is unstable 
and decomposes to manganese dioxide and permanganic acid, 
HMnO,. Permanganic acid is a comparatively strong acid and 
will at once react with the alkali present, forming a perman- 
ganate. 

Potassium Permanganate, KMn0O,, is a dark red, moderately 
soluble salt, which dissolves in water to an intensely colored red 
solution. The color is so deep that the presence of the perman- 
ganate ion is evident even in exceedingly dilute solutions. By 
reduction in acid solutions the manganese is easily reduced to a 
manganous salt, such as MnSOu,, which gives a practically color- 
less solution. In alkaline solutions the manganese is reduced to 
manganese dioxide, which is brown, leaving the solution above 
colorless. As the loss of color gives a very sharp indication of the 
end of the reaction, these properties are extensively used in 
quantitative analysis. 

As typical reactions of this character may be mentioned the 
oxidation of ferrous sulfate, FeSO,, to ferric sulfate, Fes(SO.)s, 
in the presence of sulfuric acid; the oxidation of sulfurous acid, 
H.SO;, to sulfuric acid; of nitrous acid, HNOs, to nitric acid, 
HNO;; of oxalic acid, H2C2O., to carbon dioxide, CO2; and of 
hydrogen peroxide, HO», to water, H.O, and oxygen, Ox. 

As reactions in alkaline solutions may be given the oxidation 
of sodium sulfite, Na2SO3, to sodium sulfate, Na2SOu, and that 
of manganous sulfate, MnSQ., to manganese dioxide, MnQOgz. 
The student is advised to write the equations for all of these 
reactions. 


538 A TEXTBOOK OF CHEMISTRY 


Manganese Heptoxide, or Permanganic Anhydride, Mn.O,, 
is a dark greenish black, oily, volatile liquid formed when potas- 
sium permanganate is treated with cold, concentrated sulfuric 
acid. It is extremely unstable, exploding violently on slight 
provocation. Many organic substances take fire when brought 
into contact with it. 

It will be seen on looking back through the chapter that there 
are five well defined oxides of manganese, MnO, Mn;O.1, Mn.Os, 
MnO, and Mn.O,. There is some evidence, also, but not 
entirely satisfactory, of the existence of a trioxide, MnQ3. 


CHAPTER XXXII 
IRON, COBALT, NICKEL 


Group VIII. Between manganese and copper, molybdenum 
and silver, and between tungsten and gold there is, in each case, 
a group of three metals, closely resembling each other in proper- 
ties. These three groups are: the iron group — iron, cobalt 
and nickel ; the ruthenium group—ruthenium, rhodium and pal- 
ladium ; and the platinum group — osmium, iridium and plati- 
num. The valence of these elements is. variable, as with the 
other elements of this part of the table. Valences of two, three 
and four are most common, but nickel carbonyl, Ni(CO),, ruthe- 
nium oxide, RuQ,, and osmium oxide, OsOu., indicate a maximum 
valence of eight in some cases. 

Iron, Fe, 55.84. Iron is by far the most important of the 
metals — more important than all of the other metals taken 
together. Several factors contribute to this importance. With 
the exception of aluminium, iron is more abundant than any 
other metal. As has been pointed out, oxygen forms about one 
half of the crust of the earth and silicon one fourth. Aluminium 
forms about one fourteenth and iron one twentieth, the four 
elements comprising about seven eighths of that part of the earth 
which we can examine. In addition to this, ores of iron con- 
taining 50 to 70 per cent of the metal are abundant and can be 
reduced on a large scale at a very low cost. Although alumin- 
ium is more abundant than iron, the latter can be produced 
at a very much lower cost. Finally, the forms of commercial 
iron contain carbon and other elements which greatly change 
its character, making it possible to prepare many forms of iron 
differing in hardness, malleability, tensile strength, permeability 
to magnetism and other properties which adapt the various forms 


to special uses. 
039 


540 A TEXTBOOK OF CHEMISTRY 


Occurrence of Iron. The ores of iron which are of primary 
value for manufacturing purposes are all of them oxides or com- 
pounds which can be readily converted into oxides by heat. 
The most common and important are hematite or ferric oxide, 
Fe.03, magnetite, a ferrous-ferric oxide, Fe3;0., limonite, ferric 
hydroxide, Fe:O3.Fe.(OH)s or 2 Fe.203.38 H2O, and _ siderite 
or ferrous carbonate, FeCO3. Iron pyrites, FeSe, is of value 
primarily for the sulfur which it contains, sulfur being a much 
more expensive element than iron. The oxide of iron from the 
pyrite burners of sulfuric acid plants is occasionally used for the 
manufacture of a low grade of pig iron. The presence of small 
amounts of manganese, phosphorus and sulfur affect the quality 
of the iron and are often of great importance in determining the 
value of an iron ore. Meteorites often consist largely of metallic 
iron usually containing nickel, and dredgings from the bottom of 
the ocean show the presence of meteoric dust containing iron. 
The composition of meteorites, the presence of iron in the sun 
and density of the earth as a whole (about 5.53 as compared with 
2.7 for the portion we can examine, omitting the ocean) all 
suggest the possibility that the central portions of the earth 
may contain large amounts of metallic iron. 

Metallurgy of Iron. In prehistoric times men learned how to 
reduce iron from its ores in a forge or small open hearth, with the 
use of a bellows or other device to secure a blast. The process 
was very wasteful of fuel and ore, but gave an iron of fair quality, 
sometimes approaching the properties of steel. An apparatus 
- somewhat resembling a blast furnace seems to have been invented 
about the close of the fifteenth century, but for two and a half 
centuries the ore was reduced by means of charcoal. At one 
time this use of charcoal threatened to cause the destruction of 
the forests of England. In 1735 the reduction by means of coal 
was discovered, and the use of coal and coke gradually displaced 
the use of charcoal, though some charcoal iron is still made for 
special uses, because of its purity and freedom from sulfur. 

The more important features of a modern blast furnace for 
the manufacture of iron are shown in Fig. 106. A mixture of 


a 


IRON: BLAST FURNACE 541 


ore, fuel (usually coke) and limestone is introduced at the top 
of the furnace in such proportions that the oxides of iron are 
completely reduced, giving metallic iron. The iron combines 


with carbon, silicon and 
small amounts of phosphorus 
and sulfur to form the crude 
product known as pig iron. 
The lime of the limestone 
combines with the silica and 
other impurities of the ore, 
forming a fusible silicate, 
called a slag, which melts 
and collects in the bottom 
of the furnace on top of the 
melted iron. The air for the 
combustion of the fuel is 
forced in by means of a 
powerful blower through the 
openings near the bottom of 
the furnace, which are called 
tuyéres. The fuel remains 
in excess down to the very 
bottom of the furnace, with 
the result that carbon diox- 
ide and water, formed as the 
ore is reduced, are continu- 
ally reduced back to carbon 
monoxide and _ hydrogen. 
The reduction is brought 
about chiefly by the carbon 
monoxide and hydrogen in 
accordance with the re- 
versible reactions : 


2 fou : ~ 
So mE einibiee 7 
eee a CRY pe 7 TER 
AO se PS Carer, ON gprs fen ee nN Soe ms. Paw 
= we. 2°. - Ceohey fcr “ee 7 
ss . «tas . . . 
@ ; Ce ck See ee Pei SHS one ea 
eR eee . . . 
5 is pee aieis 
tale 
df sees © 
T) Sas 
aS 
5 





an 


Fig. 106 


Peso ? het 2 CU, 
and Fe.03 + 3 Hea 2 Fe + 3 H,O 


542 A TEXTBOOK OF CHEMISTRY 


The equilibrium of these reactions is very far toward the left, 
and it is only because the solid carbon of the fuel continually 
reduces the carbon dioxide and water back to carbon monoxide 
and hydrogen that the process can succeed. These conditions 
make it necessary to use such a proportion of fuel that the gases 
escaping from the top of the furnace still contain considerable 
amounts of carbon monoxide and hydrogen — enough so that 
these blast furnace gases furnish a valuable fuel, retaining ap- 
proximately one half of the original energy of the coke. The 
gases are used to furnish the power for the blowing engines, 
either by burning them under boilers used to furnish steam for 
engines or by utilizing them directly in gas engines. The gases 
are also used in “ stoves ” in which the blast of air for the furnace 
is heated to about 800° before it enters the tuyéres. The use of 
the hot blast concentrates the reactions of the furnace in the lower 
part and greatly lessens the amount of fuel required in the charge. 
It has been discovered rather recently that a further saving of 
about 10 per cent in the amount of coke required per ton of 
iron is effected by first cooling the air for the blast to a low tem- 
perature so as to condense most of the moisture which it contains 
(Gayley; see Journal of Industrial and Engineering Chemistry, 
5, 241 (1913)). | 

The materials used in a furnace working with an ore containing 
60 per cent of iron are, approximately, in the proportion, one ton 
of ore, 0.6 ton of coke and 0.3 ton of limestone, but the amounts — 
vary with the character of the gangue in the ore and the ash of 
the coke. Approximately five tons of air must be blown into 
the furnace for each ton of ore reduced. The process is carried 
on continuously for many months and sometimes for several 
years. At intervals of a few hours the melted iron which col- 
lects in the hearth of the furnace is drawn off through the tap- 
hole at the bottom into a large ladle, by means of which it is 
transferred to a mixer, where the iron from several furnaces is 
brought together and mixed before treating it further in Besse- 
mer converters or open hearth furnaces. The larger part of the 
iron is not allowed to cool until it is converted into steel rails, 


CAST IRON 543 


steel or iron plates or structural materials of various forms. 
In the older practice, in the production of pig iron for foundry 
use or for the market, the iron was drawn out into a series of 
channels in the sand floor of the furnace room and allowed to 
solidify. It is then called pig iron. In modern practice pig 
iron is usually cast for the market in continuous casting machines. 

The slag either runs continuously from an opening above the 
surface of the iron in the hearth or more often is drawn off each 
time after the iron. The furnace slag is now extensively used 
for the manufacture of Portland cement and has sometimes been 
used to make cheap glass. | 

Pig Iron. Cast Iron. The iron from the blast furnace is 
always a crude product containing manganese, carbon, partly 
combined with the iron, partly as graphite, silicon, sulfur and 
phosphorus. ‘The larger portion of this crude iron is subjected 
to various methods of treatment which convert it into steel or 
refined irons which are more suitable for most purposes than 
the crude iron. A large amount of pig iron, however, is melted 
in cupola furnaces with or without the addition of aluminium, 
ferromanganese or other substances to improve its quality, and 
cast in sand moulds. Iron prepared is this way is gray in struc- 
ture. In the melted iron the carbon is probably all combined 
with the iron as iron carbide, Fe3C, which dissolves in the 
molten mass, forming a homogeneous solution. When the iron 
is cast as described, however, the larger part of the carbon sepa- 
rates from the iron as graphite during the slow cooling. Such 
an iron has a gray color and is known as gray cast iron. The 
presence of the graphite can be easily shown by dissolving the 
iron in a dilute acid. 

If the iron is cast in contact with a cold metallic surface, called 
a “ chill,” the carbon does not have time to separate as graphite, 
but remains combined with the iron, giving a very hard white 
iron, suitable for the rims of car wheels and other similar pur- 
poses. The addition of ferromanganese aids in holding the 
carbon in the combined form. On the other hand, silicon tends 
to cause the carbon to separate as graphite. 


544. A TEXTBOOK OF CHEMISTRY 


The following analyses indicate the usual composition of the 
material :1 


ANALYSES OF Pig IRON 











1 2 

Tron (by, difierence ia ue oan 94.29 92.72 
Conmibined Garvony sa aaa sear e. 0.55 0.61 
Grapiite Ts ae same water bees 2.22 1.85 
Len St. ORO GRE fe CAEN LON 1.84 2.57 
ELAM Sara mer os tee my REST A ee 2 5 0.035 0.044 
Ehosphorus { 4.h-ae ct eel ees 0.19 0.54 
CUUCHIAIUT sewn’. Pea a 0.074 0.081 
Manganese ole." ues an een aes 0.74 1.54 
Copper tee er oe. ae eee 0.06 0.043 

100.00 100.00 








Gray cast iron melts at 1120° to 1230°. 

Wrought Iron. While castings of gray iron can be finished 
by filing or turning in a lathe, the metal cannot be welded or 
rolled into bars or sheets. The puddling process for producing 
a nearly pure iron from the cast iron was invented by Henry 
Cort in England in 1784. The iron is melted and subjected to 
an oxidizing flame on a hearth lined with iron ore. It oxidizes 
to the magnetic oxide, Fe;0., on the surface, and by stirring 
this through the mass of the iron with a rabble the oxygen of 
the oxide burns the carbon to carbon monoxide, CO, which 
escapes. The silicon is burned to silicon dioxide, which com- 
bines with ferrous oxide, FeO, forming a fusible ferrous silicate, 
Fe:SiOy. The phosphorus is burned to phosphorus pentoxide, 
P.O;, which combines with more ferrous oxide to ferrous phos- 
phate, Fe3(PO,)2, and the sulfur burns to sulfur dioxide, which 
escapes. The process gives a malleable, ductile iron which 


1 Samples of pig iron furnished by the U. S. Bureau of Standards. 


The analyses are published with the permission of the Director of 


the Bureau. 


STEEL 545 


may be more than 99 per cent pure iron. This process, which 
was a very important one till near the close of the nineteenth 
century, has now been largely replaced by the Bessemer and 
Open Hearth processes, which are used to make mild irons as 
well as steel. Pure iron melts at 1530° and has a specific gravity 
of 7.55. It boils at 2950°. 

Cementation Steel. Cast Steel. Wrought iron contains only 
a very small per cent of carbon. It is soft and malleable and 
can be welded, but will not harden and is not suitable for the 
manufacture of knives and edge tools. By packing bars of 
wrought iron in charcoal, in long, earthenware boxes and heating 
them to 1000°-1100° for 8 to 10 days the iron absorbs approxi- 
mately one per cent of carbon and is changed to steel. The 
process is called cementation. ‘The product is melted to render 
it homogeneous and is then known as cast steel. This process, 
which was practically the only method known for making a 
good grade of steel before 1855, has been more and more dis- 
placed by the Bessemer and open hearth processes. The cemen- 
tation process is tedious and very expensive, and is now used 
only for the manufacture of a very high grade of steel for mak- 
ing cutlery and for other uses where the articles manufactured 
are small and the labor expended in giving them their final form 
is the chief item in the cost of production. The name “ cast 
steel’ is now often used for steels made by other methods. 
_ When steel is heated to a temperature of bright redness, 
700°-800°, and then cooled suddenly by quenching in water, 
it is rendered very hard. If such a steel is heated again to a 
temperature of 450°-600° the steel becomes less hard and less 
brittle. This process is called tempering. By polishing the 
steel before the second heating, the color of the film of oxide 
which forms on the surface furnishes an indication as to when 
the proper temper has been obtained. Skilled workmen deter- 
mine the temper by watching the color of the surface of the 
steel. 

*The tempering of steel can be best understood in the light 
-of the following facts: 


546 A TEXTBOOK OF CHEMISTRY 


1. Pure iron exists in three allotropic modifications called 
a-, B- and y- ferrite. The word ferrite is simply a name used by 
metallurgists to designate pure iron. These forms of iron 
differ in properties, especially in being magnetic or nonmagnetic 
and in the amount of iron carbide, FesC, which they can hold 
in solid solution. Each is stable through a definite range of 
temperature. The ranges of stability and properties will be 
seen from the following table: 














ALLOTROPIC | TEMPERATURE MAGNETIC H D z SOLUBILITY 
Form OF STABILITY PROPERTIES mma aie PSE Pas Te, or Fe3C 
a-Ferrite . . Below 750° Magnetic Soft Ductile Little 
B-Ferrite . . 750°-860° Nonmagnetic Hard Brittle Little 
y-Ferrite . . | Above 860° Nonmagnetic Hard Ductile Considerable 











The existence of these three forms has been demonstrated by 
a study of the rate of cooling of iron. If a piece of iron which 
has been heated to a temperature above 860° is connected with 
a thermocouple to record the temperature, it will be found that, 
instead of a regular fall in temperature, as would be expected, 
when a transition point is reached the temperature falls more 
slowly, ceases to fall, or may even rise for a short time, because 
of the heat evolved when y-ferrite changes to 8-ferrite or when 
8-ferrite changes to «-ferrite. In the case of steel a visible 
brightening of the mass, which is called recalescence, can be 
observed when one of the transition points is reached. 

2. Iron combines with carbon to form a definite compound, 
iron carbide, Fe3C, which is called by the metallurgists cement- 
ute. This carbide is soluble in y-ferrite but much less soluble 
in &-ferrite or #-ferrite. 

3. The presence of the iron carbide, Fe3C, lowers the tran- 
sition points so that in a steel containing 0.89 per cent of carbon 
this may all be held in a homogeneous solution in the y-ferrite 
at a temperature of 690°, or above. Such a saturated solution 
of cementite or iron carbide in y-ferrite is called austenite. It 
is very hard, and a hardened steel of this composition consists + 


BESSEMER STEEL 547 


entirely of austenite. The quenching of the steel in water 
carries it so quickly by the transition points that the transforma- 
tion does not occur and the metal is left in the hardened form. 

4. If such a steel is cooled slowly below the transition point 
to a-ferrite, the latter can no longer hold the cementite or 
carbide in solution, and the two separate into a mixture, which 
would consist, when the separation is complete, of about 13 
per cent of cementite and 87 per cent of «-ferrite. As the latter 
is soft and forms the larger part of the mass, it gives character 
to the whole. On the other hand, if the hardened steel is heated 
to 450°-600°, it changes slowly to the mixture of «-ferrite and 
cementite, the change being more rapid and complete at the 
higher temperature. The more completely the change occurs, 
the softer will be the steel. 

Bessemer Steel. In 1852 an American by the name of 
Kelly patented a process for purifying iron by blowing air 
throughit. Three 
years later, Besse- 
mer in England 
discovered _inde- 
pendently and pat- 
ented a similar 
process and_ suc- 
ceeded in develop- 
ing it to practical 
success. After 
some litigation 
Mr. Kelly sold out 
his interest to Bes- 
semer, and the 


























process is called by 
the latter’s name. 
The apparatus 
used is shown in 
ies bO7, pene he 
_large, cylindrical 





548 A TEXTBOOK OF CHEMISTRY 


vessel is at first turned on its side and a charge of several 
tons of pig iron introduced. A_ strong blast, which enters 
through one axis of the converter and through the tuyére 
holes at the bottom, is turned on and the converter brought 
to an upright position. The silicon of the iron is burned 
to silicon dioxide and the carbon to carbon monoxide. The 
silica combines with ferrous oxide to form a highly silicious 
slag. The heat from the combustion of the silicon raises the 
temperature of the iron so that it remains fluid even after the 
carbon and silicon are removed, although the melting point of 
the iron is considerably raised. When the carbon is gone, the 
flame suddenly drops and at this point the converter is again 
turned on its side and enough spiegeleisen added to give a steel 
containing the desired amount of carbon — 0.40 to 0.45 per 
cent for steel rails. After mixing, the steel is poured into ingot 
moulds and from these it is placed in “ soaking pits’ to come to 
a uniform temperature throughout and then taken directly to 
the rolls and rolled into rails or structural iron or other forms of 
iron, without being allowed to cool. 

In the “ acid’ Bessemer process as described, in which the 
lining of the converter is of silicious materials, the slag contains 
more than 70 per cent of silica, SiOs, and very little of the phos- 
phorus is removed. A good quality of steel can be obtained 
only when the material used is quite pure. To make use of 
less pure ores and iron, the basic or Thomas-Gilchrist process 
was designed. For this the converter is lined with calcined 
dolomite and lime is added to the charge. In the presence of 
the basic lining and lime the phosphorus is removed as calcium 
or magnesium phosphate. -The basic slag from the converter 
is valuable as a fertilizer because of the phosphorus which it 
contains. In America this process has been completely dis- 
placed by the basic open hearth process. 

Open Hearth or Siemens-Martin Process. During the last 
twenty years another process, illustrated in Figs. 108, 109 and 
110, has grown rapidly in favor and at the present time more 
steel is manufactured in the United States by the open hearth 





549 


OPEN HEARTH STEEL 


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RATES Se onett 
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140d Ul Y= I 
RSS 0d ULL LLL LEE 
= 2 Zt 4004 HOINS VONIS 


550 A TEXTBOOK OF CHEMISTRY 


than by the Bessemer process. In Fig. 108 two chambers are 
represented filled with a checkerwork of brick. ‘The producer 
gas (p. 297) and air used in the furnace pass up through one pair 
of these chambers, and after passing through the furnace room 
pass downward through a second pair of chambers, parting with 
their heat to the bricks. After an interval of twenty minutes 
or half an hour, the current is reversed, and now the gas and 
air take up heat from the bricks before entering the furnace 
proper. Such a furnace is called a regenerative furnace and 
works economically with a low grade of gas. In the furnace 
there is melted a mixture of cast iron, ore, steel scraps and 
sometimes lime or other fluxes. A charge of 50 to 75 tons may 
be used, and as the materials can be kept melted indefinitely, it 
is possible to take out a sample for analysis and secure a more 
accurate control than with the Bessemer converter. The 
process can also utilize an almost unlimited variety of material, 
and, especially, it is suitable for iron containing too much phos- 
phorus for the acid Bessemer and too little for the basic Besse- 
mer process. It is possible to stop the decarburization of the 
iron at any point desired instead of carrying it to completion 
and recarburizing, as is done in the Bessemer process. By all 
three processes grades of iron and steel containing from 0.1 to 
1.0 per cent of carbon are made. 

Nails and sheet iron made from materials produced by these 
modern processes corrode, when subjected to moisture and air, 
very much more rapidly than when made from wrought iron 
from the puddling process. This seems to be because the latter 
is more homogeneous and does not give differences of potential 
between different parts of the iron. It has recently been found 
that the addition of a very small amount of copper (0.2 per cent, 
or less; Chamberlain J. Ind. and Eng. Chem. 4, 360 (1913)) 
renders the iron very much more resistant to corrosion. The 
effect is similar to that produced by amalgamating the surface 
of the zinc used in an electrical cell (p. 481). 

The following analyses illustrate the composition of iron and 
steel made by modern processes. 


551 


IRON AND STEEL 


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552 A TEXTBOOK OF CHEMISTRY | 


Alloy Steels. The addition of chromium, manganese, tungsten 
and molybdenum to carbon steels lowers the point of decom- 
position of austenite (p. 546). By the use of suitable mixtures 
the point of decomposition may be brought below ordinary tem- 
perature, and such a steel will be hard at any temperature. 
Steels of this type have proved very useful for high-speed lathe 
tools, which can be run at such a rate as to become almost red- 
hot without losing their temper. Such steels are called “‘ self- 
hardening.” 

Compounds of Iron. When iron is dissolved in dilute acids, 
it forms ferrous salts, such as ferrous chloride, FeCl, or ferrous 
sulfate, FeSO, In these compounds the iron is apparently 
bivalent, but the vapor density of ferrous chloride at tempera- 
tures slightly above its boiling point points to a formula, FesCl,, 
rather than FeCl. If we assume this as the true formula, it 
still remains uncertain whether the iron or the chlorine is tri- 
valent or the iron, possibly, quadrivalent. The structure might 


be 


i] cl Cl=Cl 
Fer FeC or Fe ire 
Cl Cl Cl=Cl 


At present no means has been discovered of deciding posi- 
tively between these formulas. 

On exposure to the air, especially in neutral or alkaline solu- 
tions, ferrous compounds. are oxidized to the ferric condition. 
If an acid is present, a ferric salt is formed. If the solution is 
alkaline, neutral, or acid with a weak acid, such as carbonic 
acid, ferric hydroxide or a basic salt is formed. Iron is ap- 
parently trivalent in the ferric salts, but the vapor density indi- 
cates that ferric chloride has, in part, the formula Fe:Cl. at 
temperatures slightly above the point of sublimation. As with 
the ferrous salts, it is not known whether the molecule is held 
together by the iron or the chlorine atoms. 

The basic properties of iron are much weaker in the ferric 
than in the ferrous salts and while there is some hydrolysis in 
solutions of ferrous salts and the ferrous salts of strong acids 


IRON COMPOUNDS 553 


react acid toward litmus, the hydrolysis of ferric salts is much 
more marked. ‘This is well illustrated by the complete precipi- 
tation of the iron as ferric hydroxide when powdered barium 
carbonate, suspended in water, is added to a solution of ferric 
chloride : 

FeCl; + 3 HOH @ Fe(OH); + 3 HCl 


2 FLGT +- BaCO3 == BaCl, oe H,O -+- CO, 


The ferric hydroxide or a basic ferric salt, such as FeCl,OH 
or FeCl(OH)s, remains in collodial solution until the barium 
carbonate is added. When the barium carbonate reacts with 
the hydrochloric acid, the hydrolysis becomes complete and the 
barium chloride also assists in coagulating the colloidal ferric 
hydroxide (p. 362). 

Potassium Ferrate, K,eO., can be prepared by passing 
chlorine through a solution of potassium hydroxide in which 
ferric hydroxide is suspended. ‘This gives a red solution from 
which the salt can be crystallized. Some other ferrates have 
been prepared, but all of these hydrolyze in water and decompose 
easily. 

Ferrous chloride, FeCl:.4 H2,O, separates from concentrated 
solutions from which the air has been carefully excluded in 
clear blue crystals, which become green on exposure to the air. 
The hydrate dissolves in two thirds of its weight of water. In 
the presence of hydrochloric acid it is easily oxidized to ferric 
chloride, FeCl;, by the action of nitric acid, potassium perman- 
ganate, potassium dichromate, chlorine or almost any vigorous 
oxidizing agent. 

Ferrous Hydroxide, Fe(OH), forms as a pure white precipi- 
tate on precipitation of a ferrous salt with sodium hydroxide in 
a solution entirely free from dissolved oxygen or an oxidizing 
agent. The slightest exposure to the air causes the precipitate 
to turn green, and on longer exposure it is changed to reddish 
brown, ferric hydroxide, Fe(OH)s. 

*Ferrous Oxide, FeO, is formed as a black powder when 
ferrous oxalate, FeC2Ou,, is heated out of contact with the air. 


554 A TEXTBOOK OF CHEMISTRY 


In some conditions it takes fire spontaneously on exposure to 
the air. If reduced by hydrogen at a low temperature it forms 
pyrophorie iron, which takes fire in the air. 

Ferrous Sulfate, Green Vitriol or Copperas, FeSO,.7 H,0. 
Metallic iron dissolves readily in dilute sulfuric acid and the 
hydrate known as green vitriol crystallizes from the solution 
on cooling. One hundred parts of water dissolve 38 parts of 
the hydrate at 10° or 48 parts at 20°. Because the salt is much 
less soluble than the chloride, sulfuric acid is not so suitable as 
hydrochloric acid for the preparation of hydrogen sulfide from 
ferrous sulfide, FeS. 

Ferrous Carbonate, FeCOs, is found in nature in the mineral 
siderite. It crystallizes in rhombohedra and is isomorphous 
with calcite, CaCO3. An impure siderite, called clay iron stone, 
has been one of the most important iron ores used in England. 

Ferrous carbonate dissolves as ferrous bicarbonate, FeH2(COs3)o, 
in waters which contain carbonic acid, exactly as calcium car- 
bonate does. Such waters are known as chalybeate waters, and 
some of these have been considered valuable for their medicinal 
properties. On exposure to the air the iron of such waters is 
oxidized to the ferric state and separates as ferric hydroxide. 
Beds of iron ore were doubtless formed, in many cases by a 
similar process. 

Ferrous Chloride and Nitric Oxide. Solutions of ferrous 
chloride or ferrous sulfate absorb nitric oxide readily, giving dark 
brown or black solutions. As one molecule of the ferrous salt 
will absorb one molecule of nitric oxide, it is supposed that the 
solution contains a compound of the formula FeCl,.NO.nH,O,} 
but it has never been possible to isolate the compound, and it is 
completely decomposed by boiling the solution. 

Ferric Chloride, FeCl;.6 H.O, can be obtained by passing 
chlorine into a solution of ferrous chloride and crystallizing the 
solution. It is a yellow, deliquescent solid, very easily soluble 
in water. The anhydrous chloride, FeCl3, sublimes in dark 


1 Manchot and Zechentmayer, Liebig’s Annalen, 350, 368 (1906). 


IRON COMPOUNDS 555 


green, iridescent scales which are red by transmitted light. It 
boils, or sublimes, at 280°-285°. At 448° the weight of the gram 
molecular volume is about 303 grams and at 750° it is 167 grams. 
As the molecular weight of FeCl; is 162.2, it seems that the mole- 
cules at the lower temperature are chiefly FeCl, and at the 
higher temperature, FeCls. 

Ferric chloride is easily reduced to ferrous chloride by hydro- 
gen sulfide, by nascent hydrogen, or by stannous chloride, SnCle. 

Ferric Hydroxide, Fe(OH);, is readily formed as a reddish 
brown, flocculent, amorphous precipitate when sodium hydroxide 
or ammonia is added to a solution of a ferric salt. It does not 
dissolve in an excess of the alkali, differing very markedly from 
aluminium hydroxide, Al(OH)3, in this regard. Ferric hydroxide 
loses water very easily, and all of the hydroxides in nature con- 
tain less water than would correspond to the formula Fe(OH)3. 
The most common natural hydroxide is _ limonite, 
Fe,03.Fe2(OH)., or as it is often written, 2 Fe.03.8 H.O. Com- 
pounds of this type illustrate the somewhat artificial character 
of the distinction between hydroxides and hydrates. 

Dialyzed Iron. A solution of ferric chloride will dissolve a 
considerable quantity of ferric hydroxide. If the solution is 
dialyzed with a parchment membrane (p. 353), the hydrochloric 
acid formed by the hydrolysis of the chloride passes through 
the membrane, while the ferric hydroxide remains behind as a 
colloidal solution. Such a solution is called “‘ dialyzed iron ”’ 
and is sometimes used in medicine. It is particularly suitable 
as an antidote for arsenic poisoning. The arsenious oxide com- 
bines with the ferric hydroxide to form an insoluble arsenite. 

Ferric Oxide, Fe.O3, is prepared, artificially, by igniting the 
hydroxide. It is also found in nature as the mineral hematite. 
Ferric oxide prepared in various ways Is sold as a pigment under 
the name Venetian red or as rouge. A mixture of calcium sulfate 
and ferric oxide prepared by calcination of a mixture of sulfate 
of iron and lime is sometimes sold under the same name. The 
best quality of rouge is obtained by calcining ferrous oxalate. 
The shade obtained depends on the temperature of calcination. 


556 A TEXTBOOK OF CHEMISTRY 


*Ferric Sulfate, Fe2(SO,)3, is formed by the oxidation of 
ferrous sulfate in the presence of sulfuric acid. It forms alums, 
of which ferric ammonium alum, NH4Fe(SQO,):.12 H.O, is the 
best known. Ferric sulfate was formerly obtained by the oxi- 
dation of a pyrite-bearing shale at Nordhausen in Bohemia; 
and was decomposed by dry distillation for the preparation of 
fuming sulfuric acid. 

Magnetic Oxide of Iron, Fe;O., is the product formed when 
iron burns in air or oxygen or when steam is passed over heated 
iron. It occurs also as the natural mineral magnetite, which 
is one of the purest of the iron ores in Norway and Sweden. 
It may be considered as ferrous ferrite, Fe(FeQ2)2, or FeO.Fe.03, 
formulas which bring out its relation to chromite, with which it 
is isomorphous. 

Ferrous Sulfide, FeS, is readily formed by heating a mixture 
of iron and sulfur. It may also be prepared by heating iron 
pyrites, FeSs, which loses half of its sulfur at a high temperature. 
It is formed as a black precipitate by the addition of ammonium 
or sodium sulfide to a solution of a ferrous salt. Ferrous sulfide 
dissolves easily in dilute sulfuric or hydrochloric acid, the reaction 
which is commonly used for the preparation of hydrogen sulfide. 

* Ferric Sulfide, Fe.S3, forms as a black precipitate on adding 
ammonium sulfide or sodium sulfide to a solution of a ferric salt : 


2 FeCls +- 3(NH4)2S = Fe.S3 + 6 NH.Cl 


Ferric sulfide reacts with an ammoniacal solution of zine 
chloride to give zine sulfide and ferric hydroxide, while ferrous 
sulfide, FeS, reacts scarcely at all with the same solution. 
(Stokes, J. Am. Chem. Soc. 29, 304.) 

Iron Disulfide, or Iron Pyrites, FeSs, is a bright yellow mineral 
sometimes called “ fools’ gold ’’ because of its color and appear- 
ance. It is brittle and is easily recognized by its burning to 
sulfur dioxide and ferric oxide when heated in the air. It is 
the chief source of sulfur for the manufacture of sulfuric acid. 

Ferric Thiocyanate, Fe(CNS)3. When ammonium thio- 
cyanate, NH,CNS, is added to a solution of a ferric salt, ferric 


Sa —— a 


COBALT 557 


thiocyanate is formed and imparts a deep red color to the solu- 
tion. ‘The reaction is used for the qualitative detection and 
sometimes for the colorimetric determination of iron. 

The ferrocyanides and ferricyanides have been described in a 
previous chapter. 

Cobalt, Co, 58.94. Cobalt is usually found associated with 
nickel and combined with arsenic or sulfur or both. The best 
known mineral is smaltite (CoNi)Aso. Cobalt is a hard white 
metal, closely resembling iron. It is malleable and ductile and 
slightly magnetic. As it is not very abundant and has few 
properties which would make it distinctly more valuable than 
iron, it has not attained any considerable commercial use as a 
metal. An alloy with chromium is even more resistant to the 
attack of acids than nichrome (p. 560) and offers some promise 
of use for fruit knives and for spatulas for chemical laboratories 
(Haynes, J. Ind. and Eng. Chem. 2, 397). Cobalt melts at 1478° 
and has a specific gravity of 8.5. 

Compounds of Cobalt. Oxides. Cobalt forms cobaltous and 
cobaltic compounds, corresponding to the ferrous and ferric 
salts, but while iron tends, on the whole, to pass into the ferric 
state, the ordinary compounds of cobalt are the cobaltous salts. 
The four oxides are cobaltous oxide, CoO, cobaltic oxide, Co2Os, 
cobaltous-cobaltic oxide, Co304, and cobalt dioxide, CoQs. 

* Cobaltous Hydroxide, Co(OH).. On adding potassium hy- 
droxide to a solution of a cobaltous salt a blue, basic precipitate 
isformed. This changes on boiling to pink cobaltous hydroxide, 
Co(OH):. If cobaltous hydroxide is heated with exclusion of 
the air, it changes to light green cobaltous oxide, CoO. 

* Cobaltous Chloride, CoCl».6 HO, dissolves in water to the 
reddish or pink solutions which are characteristic of all cobaltous 
salts. The hydrated salt is also pink, but in dry air or on warm- 
ing it loses its water of hydration and changes to a deep green 
color. This is made the basis of a “ sympathetic ink.” Draw- 
ings made with a solution of cobalt chloride remain nearly in 
visible in moist air, but come out to a clear green in dry air or 
on warming gently. 


558 A TEXTBOOK OF CHEMISTRY 


* Cobalt Sulfide, CoS, is a black sulfide which dissolves only 
slowly in dilute hydrochloric acid. Nickel sulfide, NiS, conducts 
itself in the same way, and the property is often made the basis 
of a partial separation of these metals from the sulfides of iron, 
zinc and manganese, which dissolve easily and quickly in dilute 
acids. The difference seems to be due rather to the speed of 
solution than to a marked difference in the solubilities of the 
sulfides (A. A. Noyes, Bray and Spear, J. Am. Chem. Soc. 30 
483). Cobalt sulfide dissolves easily in nitric acid. 

*Cobalt Nitrate, Co(NO3)o.6 H.O, is also a pink, easily 
soluble salt. It is used in blowpipe analysis. Alumina gives a 
deep blue color when moistened with it and ignited, while zine 
oxide gives a green color. 

Cobalt Glass. Cobalt compounds impart an intense blue 
color to the borax bead, and this is used as a delicate test for the 
element. They give a similar color to glass and are used in 
the manufacture of blue glass and indecorating porcelain. Smalt 
is a deep blue silicate of potassium and cobalt which has been 
disintegrated by pouring the melted glass into water and after- 
ward ground to a fine powder. It has a variety of technical uses. 

*Potassium Cobaltocyanide, KyCoCsNe, is a salt closely 
analogous to potassium ferrocyanide, and is formed in a similar 
manner by dissolving cobaltous cyanide, CoC2Ne, in a solution 
of potassium cyanide. 

* Potassium Cobalticyanide, K;CoC,Ng, is formed by evaporat- 
ing a solution of the cobaltocyanide exposed to the air. These 
compounds correspond to the potassium ferro- and ferricyanides 
(p. 320). Nickel forms no similar compounds 

Potassium Cobaltinitrite, K;Co(NOz2)s.H.O, forms as a bright 
yellow, difficultly soluble precipitate on adding a concentrated 
solution of cobaltous acetate to potassium nitrite, or on add- 
ing a solution of sodium cobaltinitrite, Na;Co(NOz)s, to a solu- 
tion of a potassium salt. The reaction is used for the detection 
and quantitative estimation of either cobalt or potassium. For 
the determination of potassium the addition of silver nitrate, 
AgNOs, causes the formation of a still less soluble salt in which 


=P) 


COBALT. NICKEL. 559 


the potassium is partly replaced by silver, K:AgCo(NOs)¢ 
(Burgess and Kamm, J. Am. Chem. Soe. 34, 652). Nickel forms 
no similar compounds. 

Cobalt Ammines. Cobalt forms a very great number of com- 
plex compounds with ammonia. The compounds 


Luteocobalt chloride, [Co(NHs3)¢] Cl 
Purpureocobalt chloride, een aes Cl. and 
Cl 
Roseocobalt chloride, Co(NHs)5 | Cl 
H,O 


may be given as illustrations. In these compounds the cobalt 
combines with ammonia, or ammonia and water, or sometimes 
with ammonia and chlorine or other halogens or acid radicals, 
to form complex groups which, in turn, combine with acid 
radicals to form salts. Some of these compounds show forms 
of isomerism recalling that of the hydrates of the chromic 
chlorides (p. 526). A study of these compounds by Werner, 
especially, has led to the development of new views of valence 
which are important, and which differ considerably from the 
points of view developed from the study of organic compounds. 
See Werner, Z. anorg. Chem. 3, 267-330 (1893); and Neuere 
Auschauungen auf dem Gebiete der anorganischen Chemie, 
1905. 

Nickel, Ni, 58.69, is found always associated with some 
cobalt and combined with arsenic or sulfur, or both, or in the 
form of a silicate. The principal sources for the nickel of the 
world have been Sudbury in Ontario and New Caledonia, an 
island east of Australia. At Sudbury the ore is a complex sulfide 
containing copper, arsenic and other metals. The metallurgical 
process is complicated and the details are not well known out- 
side of the works where they are carried out. The first step is 
the preparation of a nickel-copper matte by smelting the ore, 
somewhat as is done with copper ores (p. 429). 

Nickel is a white metal resembling steel. It is magnetic, 
melts at 1452° and has a specific gravity of 8.8. It takes a high 


560 A TEXTBOOK OF CHEMISTRY 


polish and does not rust or tarnish so readily as steel. It is 
extensively used for plating the ornamental parts of stoves, 
handle bars of bicycles and many other articles. For this pur- 
pose it is usually deposited from a solution of nickel ammonium 
sulfate. 

Nickel is used in several alloys, the most important being 
German silver, an alloy of nickel, copper and zinc, which is 
light colored and used as a basis for silver plated ware. Ni- 
chrome, an alloy of nickel and chromium, is resistant to acids 
and is also suitable for triangles for laboratory use, for thermo- 
couples, and for electric heating devices. Other alloys highly 
resistant to acids have been prepared and there is a prospect of 
the development of important new alloys of this type. The 
five-cent coin of the United States is 75 per cent copper and 
25 per cent nickel. 

Compounds of Nickel. Nickel forms three oxides, nickelous 
oxide, NiO, nickelic oxide, NisO3, and nickelous nickelic oxide, 
Ni;04, corresponding to the similar oxides of iron. ‘The salts 
of nickel form green solutions which are complementary in color 
to the salts of cobalt. Mixtures of the two can be made which 
are nearly colorless. The chloride is NiCle.6 H2O and the 
sulfate NiSO.4.7 H.O. 

Nickel Dimethylglyoxime, 


CH ates Nae CH,— C—NOH ] 
ni , 
CH elie N—O CH,— pea a ce 


Nickel eae no precipitate with potassium nitrite, and cobalt 

can be separated from solutions containing it by this means. 
CH;—C=NOH 

It is precipitated by dimethylglyoxime, | eA , from 
CH;—C=NOH 

an ammoniacal solution or from a solution containing a weak 

acid, in the form of a scarlet-red, highly characteristic compound, 

nickel dimethylglyoximine. Cobalt forms no similar precipitate, 

especiaily if the ammoniacal solution is first shaken in the air 

to oxidize the cobalt and convert it into a complex cobaltic 


NICKEL 561 


ammine. The formation of this precipitate furnishes one of 
the best means for the detection and quantitative estimation of 
nickel. The precipitate may be sublimed without decomposition 
by careful heating (Chugaev,! Ber. 38, 2520; Z. anorg. Chem. 
46, 144; Kraut, Z. angew. Chemie, 19, 1793). 

Nickel Carbonyl, Ni(CO)s. By passing carbon monoxide 
over finely divided nickel at a temperature below 80°, nickel 
carbonyl is formed. It is a volatile liquid which boils at 43° 
and whose vapor decomposes explosively at 60°. Cobalt forms 
no similar compound, and attempts have been made to use it 
industrially for the separation of nickel from other elements, 
but these attempts have not met with much success. Iron 
forms a similar compound, iron tetracarbonyl, Fe(CO)4, and also 
a pentacarbonyl, Fe(CO);, but these are much less stable than 
the nickel carbonyl. However, it seems probable that traces 
of these iron compounds are formed from the carbon monoxide 
of illuminating gas and are the cause of the deposit of ferric 
oxide sometimes obtained from gas burners. 


1 Also written Tschugaeff. 


CHAPTER XXXIII 


THE PLATINUM METALS 


THE following table gives the atomic weights, specific gravity, 
melting points and formulas of the oxides of the metals of 
Group VIII of the Periodic System : 


























Fe Co Ni 
Atomic weight. 55.84 58.94 58.69 
Specific gravity . 7.88 8.7 8.8 
Melting point 1530° 1478° 1452° 
| FeO CoO NiO 
° Fe304 Co304 Ni304 
Sens | Fe.0; CO NiO; 
CoO, 
Ru Rh Pd 
Atomic weight . 101.7 102.9 106.7 
Specific gravity . 122 12.6 11.9 
Melting point 2300° 1940° 1549° 
RuO RhO PdO 
- Ru.O3 Rh,O3 
Sea RuOs RhO; PdO; 
| RuO* 
Os Ir Pt 
Atomic weight . 190.8 193.1 195.2 
Specific gravity . 22.48 22.4 21.4 
Melting point. 2700° 2300° 1755° 
OsO es 
5 Os.Os3 Tr.03 PtO 
pues O30: Ir» PtO. 
OsO4 


The platinum metals are found almost exclusively in the free 
state, alloyed together in small grains or nuggets. 


forms two thirds to five sixths of the alloy. 


562 


Platinum 


THE PLATINUM METALS 563 


*Ruthenium, Ru, 101.7, is found in the mineral laurite, 
(RuOs)2S3, as well as in the natural platinum alloys and in 
osmium-iridium. Ruthenium monoxide, RuO, is obtained by 
heating a mixture of ruthenium dichloride, RuCle, and sodium 
carbonate. The sesquioxide, Ru2O3, is formed when the metal 
is heated in the air. It will be seen from this that ruthenium 
resembles iron rather than platinum. Ruthenium dioxide, 
RuOs, is prepared by heating ruthenium in a current of oxygen. 
Ruthenium tetroxide, RuQu,, is a volatile compound formed by 
fusing ruthenium with potassium hydroxide and potassium 
nitrate. Its odor resembles that of ozone. Ruthenium forms 
three chlorides, RuCl, RuCl; and RuCly, and such double 
chlorides as K4RuCls (or 4 KCl.RuCl,) and K.RuCls. The 
potassium ruthenate, K,RuOx..H2O, which gives orange-red solu- 
tions and perruthenate, KRuO,.H.O, giving a green solution, 
correspond to the manganates and permanganates. 

* Rhodium, Rh, 102.9, is much less easily attacked by acids 
and other reagents than ruthenium. It is hard and has been 
sometimes used for the tips of gold pens. Rhodium monoxide, 
RhO, is obtained by heating the hydroxide, Rh(OH)3. The 
sesquioxide, Rh»2Os, is prepared by heating the nitrate, and the 
dioxide, RhOs, by fusing rhodium with potassium hydroxide 
and nitrate. The chlorides are RhCl, and RhCl;. The double 
chloride with potassium is K,RhCl;.H2O or 2 KC].RhCl;.H.0. 


There are many complex salts, such as roseorhodium chloride, 
jan (chs and luteorhodium chloride, Rh(NH3)¢6Cls. 
lee a er 


Palladium, Pd, 106.7, is always present in the natural platinum 
alloys. It is also frequently found in metallic silver. It re- 
sembles platinum or silver in appearance and occupies a posi- 
tion somewhat between them in its properties. It can be rolled 
into foil and drawn into wire. It is soluble in nitric acid, and 
finely divided palladium will dissolve in hydrochloric acid. 

One of the most remarkable properties of palladium is its 
absorption of hydrogen. If palladium foil is heated in an at- 
mosphere of hydrogen and then allowed to cool in a current of 


564 A TEXTBOOK OF CHEMISTRY 


the gas, 100 grams of the metal will absorb about 0.64 gram of 
hydrogen or approximately 7 liters of the gas. As the volume 
of 100 grams of the metal is only 8.4 cc., it follows that the metal 
absorbs more than 800 times its volume of the gas. This 
property has been used as a convenient method of weighing 
hydrogen for the determination of its atomic weight. The 
hydrogen absorbed is in an active form. It is oxidized to 
water at once in contact with oxygen or the air. A solution of 
colloidal palladium also has a similar catalytic effect and has 
been used with hydrogen for the reduction of organic compounds. 
It has even been proposed to use this method commercially for 
the reduction of liquid or semiliquid fats containing glycerides of 
unsaturated acids to convert these into solid fats which are 
commercially much more valuable. 

* Palladium forms only two well-defined oxides, the monoxide, 
PdO, and the dioxide, PdO». Palladium dichloride, PdCl».2 H.O, 
is obtained by dissolving spongy palladium in hydrochloric acid. 
The solution is reduced by hydrogen and is sometimes used to 
absorb that gas in gas analysis. The double salt, K.PdCl., 
dissolves in water to a dark red solution. 

Palladium tetrachloride, PdCl4, is known only in solution 
and is not very stable. The double salt, KePdCl, or 
2 KCI1.PdCly, is difficultly soluble in cold water and crystallizes 
in scarlet-red octahedra. Palladium forms a series of ammines 
similar to those of cobalt and rhodium. 

* Osmium, Os, 190.8, is the heaviest substance known. The 
name is given because of the strong odor of its volatile tetroxide, 
OsO,.. It is found with platinum and iridium in the alloy called 
osmium-iridium, which is insoluble in aqua regia. From this 
alloy the osmium is obtained by heating in a current of oxygen, 
which converts the osmium into volatile osmium tetroxide, 
OsO.. The reguline metal has a bluish color somewhat resem- 
bling zine. 

Osmium gives four oxides, OsO, Os203, OsOz. and OsQ,. 
The tetroxide is a white solid which dissolves slowly in water, 
but volatilizes from the solution. The vapor has a disagree- 


THE PLATINUM METALS 565. 


able, chlorine-like odor. It attacks the eyes strongly and is 
extremely poisonous. It melts at 40° and boils at about 100°. 
It is often called an acid, but has no acid properties. The 
aqueous solution is sometimes used in histology to stain or harden 
tissues. 

The chlorides of osmium are OsCl, OsCl; and OsCly. 
The double salts with potassium are OsCl;.3 KCI.6 H.O, or 
K3OsClg.6 H2O and K2OsCly. Potassium osmate, K.OsO..2 H.O, 
crystallizes in rose-red or violet octahedra. 

*Tridium, Ir, 193.1, melts about 550° higher than platinum, 
and as it does not oxidize in the air at high temperatures, it has 
proved especially useful for some forms of chemical apparatus. 
It is also used for the tips of gold pens, because of its hardness. 
Its color is between those of silver and tin. The oxides are the 
sesquioxide, Ir,O3, and the dioxide, IrO2. Both are decomposed 
at a high temperature into iridium and oxygen. 

The chlorides are IrCl., IrClz; and IrCly. The double salts 
are KsIrCls.6 HeO and KeIrCls. The latter forms dark red 
octahedra, difficultly soluble in water. 

Platinum, Pt, 195.2, is very much the most important of the 
platinum metals. Its very high melting point (1745°) and the 
fact that it does not dissolve in nitric, hydrochloric or sulfuric 
acid make it an almost indispensable metal in the laboratory. 
The fact that its coefficient of expansion is almost the same as 
that of some kinds of glass has led to its use for the leading-in 
wires of electric light bulbs. Its use as a catalytic agent to 
cause the combination of hydrogen and oxygen, and also the 
combination of sulfur dioxide and oxygen in the contact process 
for sulfuric acid (p. 175), have been given. Platinum sponge is 
used as the filtering material in the Munroe-Neubauer crucibles. 

* Platinous Chloride, PtCls, is a greenish compound formed by 
passing chlorine over platinum at 240°-250°. It dissolves in 
hydrochloric acid, giving chloroplatinous acid, H2PtCly. The 
potassium salt, K2PtCl,, is used in photography. 

Chloroplatinic Acid, H»PtCl,, is usually prepared by dissolving 
platinum in aqua regia and evaporating the solution repeatedly 


566 A TEXTBOOK OF CHEMISTRY 


to expel the excess of nitric acid, but it is extremely difficult to 
obtain a pure product in this manner. The pure compound is 
best prepared by dissolving platinum black electrolytically in 
hydrochloric acid (Weber, J. Am. Chem. Soc. 30, 29). The 
acid is easily soluble in water, giving a yellow or reddish yellow 
solution, according to the concentration. With potassium or 
ammonium salts the solution gives yellow precipitates of potas- 
sium chloroplatinate, K2PtCls, and ammonium chloroplatinate, 
(NH4)2PtCls. These compounds are much used in analytical 
chemistry. Similar compounds, many of which, however, are 
more easily soluble, are formed with organic bases. Silver 
nitrate, AgNOs, precipitates silver chloroplatinate, AgePtCl,, 
and not silver chloride, AgCl, from a solution of chloroplatinic 
acid. 

Platinic Chloride, PtCly, is obtained by heating chloroplatinic 
acid in a current of chlorine at 360°. When dissolved in water, 
it gives the compound H»2PtCl,0.4 H,O or PtCls.5 HO. The 
first formula is justified by the fact that four molecules of water 
can be readily expelled, but the fifth cannot be removed without 
loss of chlorine. 

* Platinum Disulfide, PtSs, is precipitated on passing hydrogen 
sulfide into a solution of chloroplatinic acid. It is a black pre- 
eipitate which dissolves in ammonium sulfide as ammonium 
sulfoplatinate. 

Platinum forms a long series of complex ammines. 








INDEX 


Abscissas, axis of, 43. 

Absolute, alcohol, 325; potential of 
elements, 436; temperature, 39; 
units, 33; zero, 40. 

Abstract sciences, 4. 

Acetamide, acid in liquid ammonia, 
208. 

Acetanilide, 340. 

Acetic acid, acetyl chloride from, 
245; manufacture, properties, con- 
stituent of “ liquid smoke,” 329; 
strength illustrated, 386; titration 
of, 389; solubility of calcium phos- 
phate in, 462; insolubility of 
calcium oxalate in, 465. 

Acetone, formation, preparation, 
uses, 328; solvent for acetylene, 
293. 

Acetyl chloride, formed from acetic 
acid, 245. 

Acetylene, endothermic, 292, 294; 
formation, 292, preparation from 
calcium carbide, 293; in illu- 
minating gas, 295; light from, 
293; liquid, explosive, 294; poly- 
merization, 294; solution in ace- 
tone not explosive, uses, 294; 
tetrabromide, 293. 

Acheson, graphite in colloidal solu- 
tion, 2/1. 

Acid, definition, 45, 168; properties 
from oxygen, 23; strong, defined, 
168. 

Acids and bases, writing equations 
for reactions between, 156; degree 
of ionization, table, 383; dibasic, 


defined, 183; nomenclature of, 
123; organic, structure, 328; 
strength of, definition, 167; 
strength of, illustration, 386; 


strength of in relation to solu- 
bility of sulfides, 168;  tribasic, 
defined, 183; weak and strong, 
386. 
Acid chlorides, defined, 189. 
Acidimetry, 185. 

Acidity or alkalinity of indicators 
at change of color, table, 388. 
Actinium, 475; series of elements, 

475. 
Adiabatic cooling of air at higher 
levels, 232. 


Adjective dyes, 342. 

Adsorption, 278. 

Affinity, chemical, 29. 

Agate, 348. 

Air, absorption as mixed gas by 
water, 228;| a mixture, 228; 
amount of fresh, required per 
hour, 231; analysis of by nitric 
oxide, 230; calculated weight 
of 1 liter, 229; coefficient of ex- 
pansion, 38. 

Air, composition of by volume and 


weight, 228; demonstrated by 
Lavoisier, 19; determined by 
hydrogen, 227; determined by 
mercury, 227; determined by 
phosphorus, 227. 

Air, determination of moisture in 


by weighing, by dew point, and 
by moist bulb of thermometer, 
262,58) havetaction» «2a22 “pure 
fication before liquefaction, 234; 
sources of carbon dioxide in, 229; 
weight of gram molecular volume, 
228; weight of 1 liter, 229. 

Air-slaked lime, 453. 

Alabaster, 457. 

Alberene, 349. 

Albite, trisilicate, 356. 

Albumen, 343. 

Albumoses, formed in digestion, 344. 

Alchemists, name for silver, 444. 

Alcohol, defined, 324; manufacture, 


properties, uses, absolute, dena- 
tured, 325. 

Alcoholic beverages, composition, 
Oo: 


Aldehydes, 327. 

Alfalfa, fixation of nitrogen by, 199. 

Alizarin, manufacture, 341. 

Alkali industry, history of, 400. 

Alkali metals, general properties, 
395. 

Alkali-earth metals, general proper- 
ties, 451. 

Alkalimetry, 185. 

Alkalinity or acidity of . indicators 
at change of color, table, 388. 

Alkaloids, 342. 


Allotropic forms, definition, 98; 
of phosphorus, 241; of sulfur, 
162. 


569 


570 


Alloy steels, 552. 
Alloys, fusible, 269. 


Alum, ammonium gallium, 506; 
caesium, rubidium, 424; chrome, 
Pare 


Alums, 500, potassium, 500, ammo- 
nium, ammonium ferric, chrome, 
rubidium, 501. 

Alumina, blue 
nitrate, 558. 

Aluminium acetate, mordant, 342; 
amalgam, activity of, 497; bronze, 


color with cobalt 


from electric furnace, 495, com- 
position, 497; chloride, prepara- 
tion, anhydrous, hydrate, use, 


498; exercises, 507. 

Aluminium, history of metallurgy 
of, 391; occurrence, formation 
of shales, clays, soils, 494; metal- 
lurgy, history, 495; manufacture, 
496; properties, 497; alloys, 
thermite process, 497 ; compounds, 
498; use in cast iron, 543; use in 
metallurgy, 391. 

Aluminium hydroxide, precipitation, 
499; preparation from clay, 496; 
base and acid, 499; fluoride, prep- 
aration, 499; metachloroanti- 
monate, 268; oxide, dissolved 
by sodium pyrosulfate, 408; oxide, 
preparation for manufacture of 


aluminium, 495, 496; oxide, oc- 
currence, 494; artificial, A499 ; 
uses, 500 : solution in ‘sodium 


pyrosulfate, 500; sulfate, hydroly- 
sis, preparation, use, 499, 500. 
Amalgam, ammonium, 420; sodium, 
487. 
Amalgamated zinc, conduct toward 
acids, 481. 


Amalgamation process for silver, 
441. 

Amalgams, 486. 

Amethyst, 348. 

Amide, definition of, 206; ions in 


liquid ammonia, 208. 

Amine, definition of, properties, 205. 

Amino acids, formed in digestion, 
344. 

Ammines, cobalt, 559. 

Ammonia, ‘“‘ associated ”’ liquid, 204; 
combination with acids, 202 ; deriv- 
atives of, 205; detection with 
Nessler’s reagent, 492. 

Ammonia, determination of com- 
position by volume by action of 
chlorine on, 208; by decomposition 


and recombination with electric 
discharge, 209. 
Ammonia, deviation from Boyle’s 


law, 35; formation, 201; formed 





INDEX 


by action of zine on nitric acid, 
213; liquid, solutions in, 207. 

Ammonia, preparation by hydrolysis 
of a nitride, 202; from ammonium 
sulfate, 202; from aqua ammonia, 
LNW 

Ammonia, properties, solubility, 202; 
reaction between chlorine and, 209 ; 
synthesis of, 201. 

Ammonia soda process, 
400. 

Ammoniacal gas liquors, 202. 

Ammonio-cadmium sulfate, 492; -cu- 
pric sulfate, 434 ; -cuprous chloride, 
433; -zine sulfate, 492. 

Ammonium, in ammonium amal- 
gam, 420 : bicarbonate, 423;  bi- 
carbonate, use in ammonia soda 


discovery, 


process, 412: carbonate, com- 
mercial, preparation, composition, 
use, 423; carbonate, formation, 


hydrolysis, use, 423; chloroaurate, 
450. 

Ammonium chloride, formation from 
ammonia and hydrochloric acid, 
203; manufacture, dissociation, 
volatilization of dry without dis- 
sociation, relation to Avogadro’s 
hypothesis, 421; recovery of am- 
monia from, 413. 

Ammonium chloroplatinate, 566, 428. 

Ammonium chloroplumbate, 518; 
citrate, use in analysis of fertilizers, 
331; citrate, use in determining 
citrate-soluble phosphoric acid, 
461; ferric citrate, use with potas- 
sium ferricyanide in blue prints, 
jal. 

Ammonium hydroxide, dissociation to 
ammonia and water, 204; equili- 
bria in solutions of, 420; ioniza- 
tion, 203; structure according to 
Seas theory, 207; titration of, 

9, 

Ammonium hydrosulfide, prepara- 
tion, formation of polysulfides 
from, 421; magnesium phosphate, 
decomposition of, 252; molybdate, 
naa to determine phosphorus, 529, 
53 

Ammonium nitrate, decomposition of 
exothermic, use in explosives, 215, 
423 ; preparation of nitrous oxide 
from, 214; preparation, proper- 
ties, use, 422. 

Ammonium nitrite, properties, 423; 
oxalate, use to precipitate cal- 
cium, 330; phosphomolybdate, 
use to determine phosphoric acid, 
529, 530; polysulfides, formation, 
use, 422, 


INDEX 


Ammonium salts, theory of, 203; 
sodium hydrogen phosphate, use, 
423; sulfarsenite, 261. 

Ammonium sulfate from ammoniacal 
gas liquors, 202; preparation, use, 
422; use as fertilizer, 199. 

Ammonium sulfide, preparation, 421, 
sulfide, formation of polysulfides 
from, uses, 422;  sulfostannate, 
512; trinitride, structure, 221. 


Ammono-dimercuric iodide, use in 
testing for ammonia, 493. 
Ammono-mercuric chloride, 492; 


compounds, 492; nitrate, 492. 
Amorphous, definition, 162. 
Ampere defined, 33. 
Amphibole, metasilicate, 355. 
Amphoteric compounds, definition of, 
206, 483. 

Amy] acetate, use in lacquers, 338. 
Analysis, definition, 66; hydrogen 
sulfide basis of groups in, 166. - 
Analytical chemistry, groups of, 166. 
Andrews, critical temperature, 232. 


Anesthesia, produced by nitrous 
oxide, 215. 
Anhydrite, 457, 458, soluble, 458; 


conditions for formation of, 458, 
460; vapor pressure of systems con- 
taining, 459. 

Aniline, preparation, uses, 340. 

Animal charcoal, 278. 

Animal foods, 347. 

Anion, definition, 48. 

Anode, definition, 47, 113; 
of silver peroxide on, 443. 

Anthracene, alizarin from, 341; from 
coal tar, use, 295. 

Anthracite coal, composition, 280. 

Antifebrine, 340. 

Antifriction metals, 264. 

Antimonic acids, 267. 

Antimonious acid, 265. 


deposit 


Antimony, chlorides of, 267; hy- 
droxide, 265; oxides of, 265. 
Antimony, occurrence, preparation, 


263; properties, uses, alloys, ex- 
plosive, 264. 

Antimony oxychloride, 267; penta- 
chloride, 267; pentoxide, 265; 
pentasulfide, 268. 

Antimony tetrachloride, 267; tetra- 
chloride, endothermic, 267;  tet- 
roxide, 265; trichloride, explosive 
antimony from, 264; trichloride, 
preparation, properties, hydrolysis, 
267; trioxide, 265; trioxide, tartar 
emetic from, 266; trisulfide, 268. 

Antimonyl, 266; chloride, 267; po- 
tassium tartrate, 266; sulfate, 266. 

Antipyrine, 340. 





571 


Antitoxins, 345. 

Apatite, 153, 452. 

Apoilinaris water, 309. 

Aqua ammonia, 202, 203. 

Aqua regia, 213; use to oxidize sulfur 
of sulfides, 213. 

Aragonite, 452. 

Argentum, 11. 

Argol, 330. 

Argon, atomic weight, 236; coeffi- 
cient of expansion, 38; discovery, 
235; molecular weight, 236; per 
cent in air, 228; properties, 236. 

Argyrodite, discovery of germanium 
in, 361 

Arsenic acid, oxidizing agent, 259; 
preparation, properties, salts, 259; 
transformation in steps to arsenic 
pentasulfide, 261. 

Arsenic disulfide, 260. 

Arsenic from smelting copper ores, 
256; occurrence, 256. 

Arsenic pentasulfide, 260; pentoxide, 
259; “ poison ”’ to platinum catal- 
ysis for preparation of sulfur 
trioxide, 175; preparation, prop- 
erties, uses, 257; sulfides of, 260. 

Arsenic trioxide, conversion to col- 
loidal arsenic trisulfide, 261; tri- 
oxide, formation, properties, 258; 


trisulfide, 260;  trisulfide, col- 
loidal, 261. 
Arsenious acid, salts, 259; oxide, 


oxidation by nitrogen trichloride, 
224; oxide, standard in iodimetry, 
260. 

Arsenites, 259. 

Arsenopyrite, 256; arsenic from, 257. 

Arsine, Marsh’s test, 257; com- 
pared with ammonia, 243. 

Asbestos, diaphragm of for alkali 
manufacture, 401; metasilicate, 
355; platinized, preparation of, 62. 

Assaying, 440. 

Association, water and ammonia, 204. 

-ate, suffix, use, 47; use for salts, 124. 

Atmosphere, exercises, 239. 

Atomic theory, 14; volumes, curve 
for, 137; volumes, relation to 
periodic system, 136; weight of 
chlorine, determination of, 130. 

Atomic weights, selected by law of 
Dulong and Petit, 397; selection 
of, 16, 92; table of, 10; unit for, 68. 

Atomic numbers, 138, 476. 

Atoms, probably complex  aggre- 
gates, 138; structure of, 473. 

Atropa belladonna, atropine from, 
343. 

Atropine, 343. 

Atwater, respiration calorimeter, 313, 


572 


Auric acid, 449. 

Aurum, 11. 

Austenite, relation to tempering of 
steel, 546. 

Avogadro’s law, 89, 91; exercises, 99 ; 
related to law of Dulong and 
Petit, 397; relation to laws of 
Boyle and Charles, 94. 

Az-, prefix derived from azote, 220. 

Azo-, prefix derived from azote, 220. 

Azoimide, see hydronitric acid, 223. 

Azote, name for nitrogen, 220. 


Babbitt metal, 269. 

Bacteria, killed by radiations from 
radioactive elements, 476; re- 
moval from water, 83; nitrifying, 
199. 

Badische Anilin Soda Fabrik, syn- 
thetic ammonia, 201; manufac- 
ture of indigo, 341. 

Baeyer, synthesis of indigo, 341. 

Baker, atomic weight of tellurium, 190. 

Baking powder, acid potassium tar- 
trate in, 330. 

Baking soda, 412. 

Banca tin, 508. 

Barite, 162, 468; use, 470. 

Barium carbonate, dissociation pres- 
sure, manufacture of barium oxide 


1 


from, 468; chloride, 470; chro- 
mate, 528; exercises, 476; flame 
color, 471. 


Barium, hydroxide, properties, uses, 
470; nitrate, preparation, use, 470; 
nitrate, barium oxide from, 469; 
occurrence, compounds, 468 ; oxide, 
manufacture, 468; uses, 469. 

Barium peroxide, contrast with lead 
dioxide, 518; dissociation pres- 
sure, use to prepare oxygen, for 
hydrogen peroxide, 469; peroxide 
hydrate, 469; hydrogen peroxide 
from, 84. 

Barium silicofluoride, insoluble, 350; 
sulfate, use, constituent of litho- 
pone, 470; sulfate, solubility, 471; 
sulfide, preparation, use, 470. 

Barometer, correction of readings for 
altitude, 37; correction of readings 
for glass and brass scales, 36; 
Nee hase! of readings for latitude, 

Bases, definition, 121; derived from 
ammonia, 339; how formed from 
ammonia and amines, 206; weak 
and strong, 386; and acids, writing 
equations for reactions between, 
156. 

Basicity, defined, 183. 

Batteries electric, use of zinc in, 481. 


INDEX 


Battery galvanic, reverse of electro-: 
lytic cell, 439. 

Bausfield, value of the calorie at 
different temperatures, 33. 

Bauxite, 494. 

Baxter, ‘separation of praseodymium, 
and neodymium, 504. 

Beads, borax and sodium metaphos- 
phate with blowpipe, 304. 

Bearings in machinery, phosphor 
bronze for, 431. 

Becquerel, discovery of rays, 471. 

“ Bee hive,’’ coke ovens, 278. 

Beef, use of potassium nitrate in 
salt, 418. 

Beet sugar, 333. 

Bell metal, 509. 

Benedict, respiration calorimeter, 313. 

Benzaldehyde, 328. 

Benzene in illuminating gas, 295; 
from coal tar, 294; properties, 295; 
structure, 285. 

Benzine, 289. 

Benzoic acid, cocaine a derivative of, 
343; occurrence, manufacture, use 
as a food preservative, 331. 

Beryl, 451. 

Beryl, metasilicate, 355. 

Beryllium carbonate, 452; chloride, 
451; hydroxide, 451; nitrate, 452; 
occurrence, properties, compounds, 
451; sulfate, 451. 

Berzelius, determination of composi- 
tion of water, 69; experience with 


hydrogen selenide, 190; use of 
isomer, 511. 
Bessemer converter, 547; use in 


metallurgy of copper, 429. 

Bessemer steel, history, 547 ; acid and 
basic or Thomas-Gilchrist, 548; 
soaking pits, 548. 

Bi-bivalent salts, law of solubility 
product not general for, 378. 

Bicarbonate ion, ionization of, 310. 

Bicarbonates, formation from car- 
bonic acid, 310. 

Bimolecular reactions, 150. 

Binary compounds, nomenclature of, 
29 


Biological sciences, 4. 

Biscuit, forms for earthenware, 501. 

Bisdiazoacetic acid, use in preparing 
hydrazine, 222. 

Bismuth, alloys of, 269 ; basic nitrates 
of, 270; in crude copper, 430; melt- 
ing point lowered by pressure, 269. 

Bismuth, occurrence, properties, uses, 
268; nitrate, preparation, hy- 
drolysis, 270; oxides of, 269; oxy- 
chloride, 270; ‘‘ subnitrate,’’ 270; 
trichloride, 269; trisulfide, 270. 


INDEX 


Bismuthyl chloride, 270. 

Bituminous coal, calculation of heat 
of combustion of, 44; composition, 
280. 

Bivalent, definition, 64. 

Bae ash in Leblanc soda process, 

11; 

Black oxide of manganese, history of 
uses, 535. 

Blast, heating for blast furnace, 542; 
ary, 542. 

Blast furnace, 541; gas, 298; per- 
centage composition, 299; slag, 
use for cement, 454. 

Bleaching powder, 124; manufac- 
ture, 455; properties, uses, 456; 
use in purifying water, 83. 

Bleaching by sulfur dioxide, 173; 
with chlorine, 106. 

Blindness caused by methyl alcohol, 
325: 

Blowpipe, construction and use, 304; 
oxyhydrogen, 61. 

Blueing, 321. 

Blue-print paper, 331. 

Blue vitriol, 433. 

Bodenstein, decomposition of hydrio- 
dic acid, 148; heat of formation of 
hydriodiec acid, 153. 

Body, definition, 7. 

Bohemian glass, 467. 

Boiling point, criterion of pure sub- 
stance, 12; of solutions and os- 
motic pressure, 360. 

Boisbaudran, discovery of samarium, 
505. 

Bolivia, tin from, 508. 

Bomb calorimeter, 25. 

Bone ash, 241. 

Bone black, 278. 

Borax beads in blowpipe flame, 304; 
use in blowpiping, 366. 

Borax, occurrence in California, 365; 
glass, 366; uses, 367. 

Boric acid, in Tuscany, 365; prepa- 
ration, properties, uses, 366; 
use as food preservative, 366; tests 
for, 367, 368. 


Bornite, 428. 

Boron, exercises, 368; fluoride, 
hydrolysis, 367; nitride, 367; oc- 
currence, preparation, properties, 


365; sulfide, 367; trichloride, 367 ; 
trioxide, 365. 

Borosilicate glasses, 467. 

Boyle, law of, 34. 

Boyle’s law, deviation of gases from, 
35; illustration of, 35. 

Brandt, discovery of phosphorus, 242. 

Brandy, 325. 

Brass, 431. 


573 


Brazil, diamonds from, 275. 


Breath, source of carbon dioxide, 
229. 
Brick, 501. 


Brimstone, roll, 161. 

Brines, evaporation of, 405. 
Britannia metal, 264, 509. 

British thermal unit (B. T. U.), 26. 
Bromate, sodium, 144. 


Bromine, occurrence, 140; origin of 
name, 141; preparation, 140: 
properties, 141; uses, 142. 


Bronze age, relation to metallurgy, 
390; phosphor, 431. 

Bronzes, 431, 509; sodium tungstate, 
Bak: 

Brownian movement, 96. 

Bullion, lead, separation of silver 
from, 440. 

Bunsen, discovery of rubidium and 
cesium, 424; discovery of spec- 
trum analysis, 424. 

Bunsen burner, nature of flame, 300; 
burner, temperature of flame, 303; 
flame, separated, 301. 

Burgess, table of melting points of 
elements, 373. 

Burgess, use of potassium silver 
cobaltinitrite in analysis, 559. 

Burns produced by radioactive ele- 
ments, 476. 

Butadiene, structure, 285. 

Butane, normal, structure, 284. 

Butene, 283. 

Butine, 283. 

By-products coke ovens, 279. 

By-products, importance in Leblanc 
soda process, 401. 





Cadmium, abnormal ionization of, 491 ; 
compounds, effect of ammonium 
hydroxide on solutions of, 491; 
hydroxide, 484. 

Cadmium, occurrence, properties, 
483, use in fusible alloys, 484; 
separation from copper with po- 
tassium cyanide, 435 ; sulfate, 484; 
sulfate, use in Weston cell, 438; 
sulfide, conduct toward acids and 


potassium cyanide, 484; sulfide, 
solubility, 491. 
Cesium alum, 424; chloroiodide, 


use in purification of cesium, 424. 
Cesium chloroplatinate, 393, 424. 
Cesium, discovery, properties, 424: 

tetrachloroantimonate, 267. 
Cailletet, liquefaction of air, 233. 
Calamine, 481. 

Calamine, orthosilicate, 355. 
Calcite, 452. 


; Calcium acetate, manufacture, use, 


574 


465; 
310. 
Calcium carbide, hydrolysis to cal- 
cium hydroxide and _ acetylene, 
293; manufacture, hydrolysis, use, 
462; calcium cyanamide from, 462. 
Calcium carbonate, dissociation, phase 
rule, 453 ; dissociation pressures of, 
453; in mortar, 454; solubility in 
pure water and in water containing 
carbonic acid, hard waters, 463. 
Calcium chlorate, formation, use in 
making potassium chlorate, 456; 
chlorate, use to prepare potassium 
chlorate, 127; chloroaurate, 450. 
Calcium chloride, by-product in am- 
monia soda process, 413; effect 
on the formation of gypsum, 460; 
moisture left in gas by, 54; prep- 
aration, hydrates, by-product in 
ammonia soda process, as drying 
agent, uses, 455. 


bicarbonate, hard waters, 


Calcium cyanamide, manufacture, 
hydrolysis, 462; use, 463; exer- 
cises, 476; flame color, 471; 
fluoride, properties, uses, 456; 


fluoride, occurrence, 153; hydride, 
452; hydrosulfide, action of car- 
bonic acid on, 457; hydroxide, 
in setting of cement, 455; hy- 
pochlorite, manufacture, 455, prop- 
erties, uses, 456; manganite, 
formation in Weldon process, 103; 
metachloroantimonate, 268. 

Calcium nitrate, formation in soil, 
199; occurrence, potassium nitrate 
from, 418; preparation, manu- 
facture, uses, 460. 

Calcium nitride, formation, hydroly- 


sis, 452; occurrence, preparation, 
properties, 452; orthoplumbate, 
516. 


Calcium oxalate, precipitation, solu- 
bility in strong acids, insolubility 
in weak acids, use as test for 
calcium and oxalic acid, 465. 

Calcium oxide, manufacture, 452; 
equilibrium with carbon dioxide 
and calcium carbonate, 453. 

Calcium phosphates, occurrence, 460; 
superphosphates, 461; solubility, 
rere to fertilizing value, 461, 

Calcium phosphate, 
weak acids, 462; selection of 
atomic weight, 397; silicate, prep- 
aration, occurrence, 466; stearate, 


Sone 
457 ; 


solubility in 


Calcium sulfate, forms, uses, 
hydrates of and the phase rule, 
458; in cement, 454; in hard 








INDEX 


waters, 463; permanent hardness 
from, 311. 
Calcium, sulfite, acid, preparation, 


use in paper manufacture, 457 ; 
sulfite, acid, use in paper making, 
175% sulfid) in Leblane soda pro- 
cess, 411; sulfide, preparation, 
hydrolysis, 456, recovery of sulfur 
from by Chance process, 457; super- 
phosphate, manufacture, use, anal- 
ysis, 461. 

Calculation of formula of mineral, 
356; of relative speed of reactions 
at equilibrium, 151. 

Calomel, preparation, vapor density, 
molecular weight, uses, 489. 

Calorie, definition, 26; value at 
different temperatures, 33; value 
in joules, 33. 

Calorimeter, 25. 

Calorimeter, respiration, 313, muscu- 
lar energy in, 315; mental work in, 
316; cuts of, 314, 315. 

Camphor, in celluloid, 338. 

Candle power, standard, 293, of 
illuminating gas, 296, of acetylene, 
293, of Welsbach light, 296. 

Cane sugar, 333. 

Cannel coals, 281. 

Cafion Diablo, diamonds in meteo- 
rite from, 274. 

Caramel, 333. 

Carat, defined, 448. 

Carbohydrates, defined, 332. 

Carbolic acid, see phenol. 

Carbon, amorphous, heat of com- 
bustion, 276; preparation, prop- 
erties, 277. 

Carbon bisulfide and nitric oxide, 
flame of, 217; preparation, prop- 
erties, uses, formation of sulfo- 
carbonates from, 317. 

Carbon, chemical properties of, 281; 
cycle of in nature, 312. 

Carbon dioxide, amount formed in 
world by burning coal, 229; 
amount i in the ocean, 230; amount 
in air kept constant by. ocean, 


230; coefficient of expansion, 38; 
conditions for escape from solu- 
tions, 376. 


Carbon dioxide, critical temperature 
of, 233; density, 308; accumula- 
tion in wells and caves, 309, 
diffusion in rooms, 309; deviation 
from Boyle’s law, 35; exercises, 
322; formerly considered poison- 
ous, 231 ; from burning charcoal, 238. 

Carbon dioxide, limit for in ventilated 
rooms, 231; per cent in air, 228; 
preparation from calcium  ear- 


INDEX 


bonate, sodium bicarbonate, mag- 


nesium carbonate, 306;  proper- 
ties, isothermals, 307. 
Carbon dioxide, ratio of specific 


heats of, 237; reduction in plants, 
312; removed from air by plants, 
229; solubility, 309; sources, 313; 
sources of in air, 229. 

Carbon, effect on decomposition of 
barium carbonate, 468; electrodes, 
279; gas, 279; heat of combus- 
tion, 27 ; in steel, use of potassium 
cupric chloride in determining, 432. 

Carbon monoxide, absorbed by cu- 
prous chloride, 433; coefficient of 
expansion, 38; deviation from 
Boyle’s law, 35; effect of forma- 
tion on decomposition of barium 
carbonate, 468; formation from 
calcium oxalate, 466. 

Carbon monoxide, formation in 
burning coal, preparation from 
oxalic acid, 311; properties, failure 
to burn when dry, 311; in water 
gas, danger, 297; poisonous, 312; 
formation of sodium formate from, 
abe 

Carbon, occurrence, number of com- 
pounds, importance, 273; oxy- 
chloride, 316; oxysulfide, forma- 
tion from thiocyanates, properties, 
hydrolysis, 318; slowness of reac- 
tion, 280; suboxide, 316; tetra- 
chloride from methane, 287; va- 
lence, 282. 

Carbonate, ammonium, 423. 

Carbonate ion, réle in decomposition 
of carbonates, 375. 

Carbonates, decomposition by acids, 
theory, 375; formation from ecar- 
bonic acid, 310; oxides prepared 


from, 392. 
Carbonic acid, carbonates and bi- 
carbonates from, 310;  deter- 


mination of free and combined, 
464; formation, properties, 309. 
Carbonyl chloride, preparation, prop- 
erties, 316, hydrolysis, urea from, 
B17. 
Carborundum, 349. 
Carboxyl, characteristic group 
organic acids, 328. 
Carnallite, 414, 478; 
424. 

Carnotite, uranium in, 531. 

Caro’s acid (permonosulfuric acid), 
188. 

Casein, a protein, 343. 

Cassiterite, 508. 

Cast iron, composition, gray, white, 
chilled, 543; analyses, 544, 


of 


rubidium in, 


575 


Castner-Kellner apparatus for alkali 
manufacture, 402. 

Cast steel, 545. 

Catalysis, 28; aefinition, 62; in 
sulfuric acid manufacture by oxides 
of nitrogen, 178; of preparation 
of sulfur trioxide, 175; of synthesis 
of ammonia, 201. 

Catalyzer, copper chloride for Deacon 
process, 103. 

Cathode, definition, 47, 113. 

Cation, definition, 48. 

Cavendish, analysis of air by nitric 
oxide, 230; nearly discovered 
argon, 235. 

Celestite, 468. 

Celluloid, 338. 

Cellulose, use as fuel, as food, in 


paper, 337. 
Cement, dental, composition, 482; 
manufacture, composition, 454; 


setting, 455. 

Cementation steel, 545. 

Cementite, iron carbide, 546. 

Centimeter-gram-second system, 33. 

Cerargyrite, 439. 

Cereal, composition of, 8. 

Cerium, alloy with iron, 364; group 
of rare earths, 503; occurrence, 
properties, 362; oxides, sulfate, 
double sulfate with sodium, 364; 
phosphate, 363. 

Chalcedony, 348. 

Chalcocite, 428. 

Chalcopyrite, 428. 

Chamberlain, use of copper to pre- 
vent corrosion of iron, 550. 

‘* Chamber process ’”’ for sulfuric acid, 
Lie 

Chameleon solution, 536. 

Chance process for recovering sulfur, 
457. 

Chapin, separation of praseodymium 
and neodymium, 504. 

Charcoal, animal, 278; burning in 
oxygen, 23; composition, 280; 
manufacture, 277 ; properties, uses, 
adsorption by, 278. 

Charles, law of, 38. 

Chemical action of radioactive rays, 
475. 

Chemical activity in solutions, 81. 
Chemical affinity, 29; relation to 
speed of chemical reactions, 149. 
Chemical energy, nature of, 27; 

defined, 34. 

Chemical reactions, equilibrium in, 
108; effect of concentration on, 
24; speed of, 149; unimolecular 
and bimolecular, 150. 

Chemistry, definition, 5; study of, 18. 


576 INDEX 
Chicago, typhoid fever in, from |Chromic anhydride, preparation, use 
water, 83 to oxidize carbon, 528. 


Chili saltpeter, 210; potassium ni- 
trate from, 418. 

Chlorates, 127. 

Chloric acid, 127; structure, 130. 

Chloride, acid, hydrolysis of, 317; 
of lime, manufacture, 455; proper- 
ties, uses, 456. 

Chlorides, hydrolysis of, 115; of 
acids, 189; preparation by use of 
sulfur monochloride, 188. 

Chlorination process for gold, 446. 

Chlorine and oxygen, comparison of 
heats of combination, 108; action 
on ammonia, 209; bleaching by, 
106; burning in hydrogen, 118. 

Chlorine, combination with other ele- 
ments, 104 ; determinationof atomic 
weight of, 130; effect of light on 
reaction with hydrogen, 105; effect 
of moisture on combination with 
other elements, 105; exercises, 116; 
from electrolysis in alkali manufac- 
ture, 402. 

Chlorine hydrate, phases, 107; list 
of oxides and oxygen acids of, 123; 
occurrence, 100; dioxide, 127. 

Chlorine, preparation by electrolysis 
of sodium chloride, 100; by the 
Deacon process, 102; from hydro- 
chloric acid and manganese dioxide, 
101; from hydrochloric acid and 
potassium permanganate, 102; by 
oxidation of hydrochloric acid, 100 ; 
by Weldon process, 102. 

Chlorine, properties, 104; 
with water, 106. é 

Chlorites, 127. 

Chloroamine, 224. 

Chloroaurates, 450. 

Chloroauric acid, 450. 

Chlorocuprous acid, 432. 

Chloroform from methane, 287. 

Chloroplatinates from amines, 423. 

Chloroplatinic acid, preparation, 565; 
from platinum black, 566; use in 
photography, 445. 

Chloroplatinous acid, 565. 

Chloroplumbic acid, 518. 

Chloroplumbous acid, 518. 

Chlorosulfonic acid, preparation, 
properties, 189. 

Chlorous acid, 127; structure, 130. 

Cholera from impure water supply, 83. 

Chrome alum, preparation, 527. 

Chrome green, 525. 

Chrome iron ore, 524; decomposition 

Chrome tanning, 527. fof, 527. 

Chrome yellow, 527; constituent of 
chrome green, 525. 


reaction 





Chromic chloride, hydrates, 525; 
isomeric, 526; preparation, proper- 
ties, 525; theory of, 526. 

Chromic hydroxide, formation, com- 
position, properties, 525; chro- 
mites, 525. 

Chromic oxide, preparation, use as 
pigment, 525. 

Chromite, 524. 

Chromites, 525. 

Chromium, alloy with cobalt, 557; 
occurrence, metallurgy, properties, 
uses, 524; preparation by thermite 
process, 524, 498. 

Chromium trioxide, preparation, use 
to oxidize carbon, 528. 

Chromous' chloride, preparation, 
properties, 525. 

Chromyl chloride, preparation, prop- 
erties, structure, hydrolysis to di- 
chromic acid, 528. 

Chugaev, separation of nickel and 
cobalt, 561. 

Cinnabar, 485, 489. 

Citral, use in making ionone, 328. 

Citrate-soluble phosphoric acid, 461. 

Citric acid, structure, source, 330; 
salts, 331. 

Clark cell, electromotive force, 437. 

Clarke, F. W. Composition of the 
crust of the earth, 11. 

Classification, of metals, 370, 371; 
of the elements, 132. 

Clays, formation of, 494; manufac- 
ture of aluminium oxide, hydro- 
chloric acid and sodium carbonate 
from, with salt, 496. 

Cleveite, discovery of helium in, 237; 
uranium in, 531. 

Clinker, cement, 454. 

Clouds, conditions of formation, 232. 

Clover, fixation of nitrogen by, 199. 

Coal, energy of from sunlight, 230; 
formation, varieties, composition, 
280. 

Coal gas, percentage composition, 
299; tar, character, 294; tar dips, 
phenol in, 326. 

Coals, coking, noncoking and cannel, 
281. 

Cobalt ammines, 559; 
558; nitrate, use 
analysis, 558. 

Cobalt, occurrence, properties, alloy 
with chromium oxides, 557; sepa- 
ration from nickel by dimethyl- 
glyoxime, 560; sulfide, formation, 
slow solubility, 558. 

Cobaltous chloride, properties, sym- 


glass, smalt, 
in blowpipe 


INDEX 


pathetic ink, 557; hydroxide, prep- 
aration, properties, 557. 

Cocaine, 343; similar synthetic alka- 
loids, 348. 

Coefficients of expansion of air, On:, 
Ne, NO, He, A, He, CO, COz2, SOx, 38. 

Coining value of gold, 447. 

Coins, gold, 448; nickel five cent, 
560; silver, 442. 

Coke, composition, 280; 
ture, uses, 278. 

Coking coals, 281. 

Colemanite, boric acid from, 365. 

Collodion, 338. 
Collection and storage of gases, 22. 
Colloidal arsenic trisulfide, 261; 
silicic acid, 353; solutions, 262. 
Colloids, contrasted with crystalloids, 
357; precipitated by bivalent ions, 
263; properties, 262. 

Columbite, 523. 

Columbium (niobium), discovery, oc- 


manufac- 


currence, properties, compounds, 
B23: 
Combining volumes, law of, 89; 


weights, law of, 13. 

Combustion, 24; heat of, 25. 

Complex cyanides, 319. 

Complex ions, formation of, 378; 
evidence for existence of, 379. 
Composition of air demonstrated by 

Lavoisier, 19; of the crust of the 
earth, 11; of pure substances 
expressed in multiples of atomic 

weights, 17. 

Compounds, definition, 9; general 
methods of preparing, 372-379. 
Compressed gases, cooling on ex- 

pansion, 233. 

Comstock, dependence of mass on 
velocity, 5. 

Concentration and speed of reaction, 
149; effect of on chemical reac- 
tions, 24. 

Congress water, 309. 

Coniine, 342. 

Conservation of energy, 6; of matter, 


Constant proportion, law of, 12. 

“Contact mass” for catalyzing 
formation of sulfur trioxide, 176. 

Converter, Bessemer, 547. 

Cooking of starchy foods, 336. 

Copper, addition to steel, 431; action 
of nitric acid on, 213; acetylide, 
see copper carbide. 

Copper, alloys of, 481; annual pro- 
duction, value, 4380; properties, 
effect of impurities on conduc- 
tance of, 430; uses, 431; arsenite 
and acetate (Paris green), 259. 








577 


Copper, basic carbonate of, 
formation, 431. 

Copper carbide, preparation, 292; 
chloride, catalyzer for Deacon 
process, 103; detected by sodium 
metaphosphate, 253; electro- 
lytic refining, 429. 

Copper, exercises, 450; ferrocyanide, 
use in semipermeable membranes, 
358; hydroxide, precipitation, de- 
composition, 431; in five cent 
piecé, 560; metallurgy of, 428. 

Copper, occurrence, metallurgy, 428; 
electrolytic refining, 429. 

Copper oxide, from copper hydroxide, 
431; from wire, nitrate, use, 432; 
[we LOXONG uae? lo nse sin wdeter= 
mination of the composition of 
water, 69. 

Copper, precipitation by 
thiosulfate, 409; precipitation by 
iron, 435; prevention of corro- 
sion of iron by, 550; pyrites, 161, 
428; separation from cadium as 
cuprocyanide, 435. 

Copper, sulfate, hydrates, uses, 433; 
reaction with sulfuric acid, 173; 
titration of with potassium iodide 
and sodium thiosulfate, 433. 

Copperas, 554. 

Corn sirup, 334. 

Cornwall, tin from, 508. 

Corpuscle, same as electron, 181. 

Correction of volume of gas for 
temperature, 39; zero and stem, 
for thermometers, 486. 

Corrections for readings of barom- 
eter for altitude, 37; for read- 
ings of barometer for glass and 
brass scales, 36; for readings of 
barometer for latitude, 37. 


428; 


sodium 


Corrosion of iron, prevention by 
copper, 550. 

Corrosive sublimate, preparation, 
489; properties, uses, antidote 
for, 490. 


Cort, invention of puddling process, 
544 


Cotton goods, fire-proofing of, 513. 

Cowles Brothers, electric furnace, 
495; use of electric furnace for 
aluminium, 391. 

Cowles, process for decomposition 
of clay, 496. 

Cowper-Cowles, Sherard, 
ized iron, 482. 

Cream of tartar, tartar emetic from, 
266; use in jellies, in baking pow- 
ders, 330. 

Cretinism, connected with  defi- 
ciency of iodine, 144. 


sherard- 


578 


Critical temperature, defined, 233; 
discovery of, 232; of carbon 
dioxide, 233; relation to lique- 
faction of air, 233. 

Crookes, discharge of electricity 
through rarefied gases, 471; dis- 
covery of thallium, 507. 

Crown glass, 467. 

Cryolite, 153, 495; use for manu- 
facture of aluminium, 495. 

Crystallization, 8; water of, 82. 

Crystallographic systems, 193. 

Crystalloids, 357. 

Crystals, definition of, 192. 

Cubic centimeter, true volume of, 31. 

Cupellation, 440. 

Cupric chloride, preparation, ioniza- 
tion, 432; nitrate, hydrates, de- 
composition of, 434; oxide, prep- 
aration, use, 432; sulfide, forma- 
tion, properties, 433. 

Cuprous chloride, preparation, prop- 
erties, 432; use to absorb car- 
bon monoxide, 433; cyanide, for- 
mation, complex salt with potas- 
sium cyanide, 434; iodide, forma- 
tion in titrating copper, 433; 
oxide, formation, in testing for 
glucose, hydrazine, etc., proper- 
ties, 432; sulfide in matte, 429; 
sulfide, occurrence, 428; forma- 
tion, 433. 

Cyanides, complex, 319; formation, 
preparation, 319; from ammoni- 
acal gas liquors, 319. 

Cyanide process, for gold, 446; 
for silver, 441. 

Cyclopentene, structure, 285. 

Cyclopropane, structure, 284. 


formula 

law of multiple 
proportion, 88; law of partial 
pressures, 41, 77; view of mole- 
cules of the elements, 93. 

Damascus blade, tungsten in, 530. 

Davy, discovery of metallic potassium, 
415; discovery of metallic sodium 
and potassium, 399; injured by 
nitrogen trichloride, 224; safety 
lamp, 287; study of explosion of 
fire damp, 287. 

Deacon process for chlorine, 102; 
equilibrium of, 109. 

Debye, quantum theory, 398. 

Degree of ionization, measurement of, 
380; table, acids, 383 ; bases, 383 ; 
salts, 384. 

Degrees of freedom, 77. 

Dehydration by sulfuric acid, 182. 

Dekahydronaphthalene, 285, 


Dalton, atomic theory, 14; 
for water, 91; 





IN DEX 


Deliquescence, 82. 

Density, criterion of pure substance, 
12; of gases, table, 95. 

Dental cement, composition, 482. 

Derivatives of ammonia, 205. 

Destructive Distillation, 328. 

Determination of weight of a liter of 
gas, 40. 

Detonating caps, 491. 

Developing in photography, 445. 

Dewar flasks for liquid air, 235. 

Dextrin, manufacture from starch, 

Dextrose, see Glucose. [uses, 336. 

Dialysis, 357. 

Dialyzed iron, 555. 

Diamond, heat of combustion, 276; 
artificial, 275, natural, 275; uses, 
properties, 275. 

Diastase, 344. 

Dibasic acids, defined, 183. 

Dibromoethane, 291. 

Dicalcium phosphate, 249. 

Dichloroethane, 291. 

Dichromic acid, from hydrolysis of 
chromyl chloride, 528. 

Diet, salt essential in, 406. 

Dietary, average American, 347. 

Diffusion of gases, 56; law of, 59. 

Digestion, colloidal solutions, 263. 

Dimethylglyoxime, precipitant for 
nickel, 560. 

Diphosphorus pentasulfide, 254. 

Disilicates, 356. 

Disilicic acid, 355. 

Disinfectant, formaldehyde, 
sulfur dioxide, 174. 

Disintegration of atoms, possible by 
radium emanation, 475. 

Disodium phosphate, 249; hydroly- 
sis, alkaline reaction of, 251; 
phenolphthalein as indicator for, 
251; uses, arsenic as impurity in, 
410. 

Displacement of equilibrium, 152. 

Dissociation, definition, 59; pressures 
of calcium carbonate, 453; pres- 
sure of silver oxide, 443; of am- 
monium hydroxide to ammonia 
and water, 204; of calcium car- 
bonate, phase rule, 453; of sulfuric 
acid, 180; of water, 59. 

Distillation, 8. 

Dithionic acid, 188. 

Divariant, definition, 77. 

Dixon, explosion waves, 301. 

Dolomite, 478. 

Double decomposition, defined, 81. 

Double refraction of crystals, 196. 

Dry batteries, use of zinc in, 482. 

Dry plates, photographic, 444. 

Drying of gases, 54, 


BPA 


ee ~~ 


INDEX 


Dulong and Petit, law of, 396; 
relation to Avogadro’s law, 397. 
Dumas, determination of the compo- 

sition of water, 69. 

Durax glass, 467. 

Dust, explosion of with air, 289. 

Dutch process for white lead, 520. 

Dydimium, separation into praseo- 
dymium and neodymium, 504. 

Dyeing, lakes for, 501. 

Dyes, 340; use of sodium nitrite in 
manufacture of, 410; substantive, 
adjective, 342. 

Dysprosium, compounds, 506. 


Earth, composition of crust of, 11; 
mean density of, 540. 

Earthenware, 501, glazing, 502. 

Effervescent waters, 309. 

Efflorescence, 82. 

‘Egypt, early manufacture of iron in, 
390; sodium carbonate from, 450. 

Einstein, quantum theory, 398. 


Eka-aluminium, same as gallium, 506. 


Ekaboron, same as scandium, 136, 503. 
Ekeberg, discovery of tantalite, 523. 
Electric furnace, carborundum in, 
349; use for aluminium, 391. 
Electrical batteries, use of zinc in, 
481, theory of, 435; horse power, 
34; unit charge, 438; units, 33. 
Electrochemical theory, influence on 
formulas of minerals, 356. 
Electrodes, carbon, 279. 
Electrolysis, migration of ions in, 113; 
of dilute sulfuric acid, 47, 9; 
sodium hydroxide by, 402. 
Electrolyte, definition, 48. 
EES methods .in metallurgy, 
i 
Electromotive force, relation to solu- 
tion pressure, 435. 
Electromotive series, 435, table, 436. 
ee relation to Faraday’s law, 
Electron theory, 181; as explanation 
of ionization, 182; relation to 
ionization, 206; relation to prop- 
erties of metals and non-metals, 


370. 

Sie creesnye elements, defined, 
437. 

Electroplating, copper, 433. 

pec eoponive elements, defined, 
37. 


Electrotyping, 433. 

Elements, absolute potential of, 436; 
atoms of probably complex aggre- 
gates, 138; classification of, 132; 
definition, 9; life of, 474; melting 
points of, absolute, 135; melting 





579 


points of, 372, table, 373; metallic 
in periodic system, 136; missing 
in Group VII, possible reason, 5338 ; 
molecules of, 93. 


-Elements, non-metallic in periodic 


system, 136; radioactive, series of, 
475; specific heat of, 397; symbols 
of, 11: table; .of (familiar,..371; 
table of groups of, 371; table 
of non-metallic, 348. 

Emery, 494; artificial, 500. 

Endothermic compounds, defined, 
225; explosive decomposition, 225. 

Endothermic reactions, 215. 

Energy, conservation of, 6; defini- 
tion, 6; muscular, in respiration 
calorimeter, 315; of coal from 
sunlight, 230. 

Energy required to decompose a 
gram equivalent, 438; units of, 32; 
varieties of, 6. 

England, destruction of forests for 
iron manufacture, 540. 

English laws for glazes, 502. 

Enzymes, 344. 

Epsom salts, 478, 480. 

Equations, writing of, 49; writing of, 
for reactions between acids and 
bases, 156. 

Equilibrium between gaseous and 
solid phases, 443, silver, silver 
oxide and oxygen, 4438; between 
water and water vapor, 76; dis- 
placement of, 152; effect on, of 
removing one of the reacting sub- 
stances, 152; for combination of 
nitrogen and hydrogen, 201; for 
formation of nitric oxide, 216; 
hydrogen, iodine and hydriodic 
acid, 146. 

Equilibrium in chemical reactions, 
108; in gas flame, 300; illustra- 
tion of that between water and 
water vapor, 76; in ionization of 
orthophosphoric acid, 250; in 
neutralization, 385. 

Equilibrium of carbon monoxide, 
carbon dioxide, hydrogen and water 
vapor, 300; of Deacon process, 
109; of reaction between hydro- 
chloric acid and oxygen, 108. 

Equivalents, relation to Faraday’s 
law, 4388. 

Erbium, compounds, 506. 

Erg, defined, 32. 


Ethane, structure, 284;  substitu- 
tion products of, 291. 
Ethene, 290. 


Ether, ethyl, 290. 
Ethyl alcohol, ether and ethylene 
from, 290; ethyl chloride from, 


580 


245; manufacture, properties, uses, 
325; absolute, denatured, 325; 
structure, 323. 

Ethyl borate, 367; chloride, formed 
from ethyl alcohol, 245; ether, 
formation, 290; iodide, relation 
to structure of ethyl alcohol, 323. 

Ethylene, addition compounds, 291; 
bromide, 291; chloride, formation 
from ethylene, 291; in illuminat- 
ing gas, 295; preparation, prop- 
erties, uses, formation and de- 
composition, 290; ratio of specific 
heats of, 237; structure, 292. 

Eudiometer, description, 67. 

Europium, compounds, 505. 

Eutectic point, defined, 488. 

Evaporators, triple and multiple- 
effect, 405; Yaryan, 405. 

Exercises, aluminium, 507; atmos- 
phere, 239; Avogadro’s law, 99; 
calcium, barium, 476; carbon 
dioxide, cyanides, 322; chlorine, 
116; copper, silver, gold, 450. 

Exercises, Group V, 272; hydrocar- 
bons, 305; laws of gases, 43; mag- 
nesium and mercury, 493; nitro- 
gen, 225; phosphorus, 255; silicon, 
boron, 368; sulfur, 196; writing 
equations, 156. 

Exothermic reaction, definition of, 
ilies 

Explosions, definition, 62; of en- 
dothermic compounds, 225; of 
methane and air, 287; of dust and 
air, 289; waves, 301. 


Factories, humidity in, 232. 

Families of elements, table, 371. 

Faraday, injured by nitrogen tri- 
chloride, 224; law, 338; liquefac- 
tion of chlorine, 108. 

Fast colors, 340. 

Fats, composition, 331; use in 
making soaps, 332; reduction with 
the aid of colloidal palladium, 564 ; 
soft soap from, 414. 

Fatty acids, calcium salts of, 332. 

Fehling’s solution, formula for, use, 
to detect glucose, 335. 

Feldspars, 348; use as glaze, 502. 

Ferric acetate, mordant, 342. 

Ferric chloride, hydrate, anhydrous, 
554; molecular weight, 555; 
hydrolysis, precipitation of ferric 
hydroxide from, 5538; theory of 
hydrolysis of, 386. 

Ferric ferrocyanide, 320; decomposi- 
tion with sodium hydroxide, 321; 
hydroxide, precipitation, 555; by 
barium carbonate, 553; hydroxide, 











INDEX 


to absorb hydrogen sulfide, 295; 
oxide, preparation, manufacture, 
use, 555. 

Ferric sulfate, fuming sulfuric acid 
from, 556; preparation, alums from, 
556; reduction by hydrogen sul- 
fide, 171; sulfide, formation, prop- 
erties, 556; thiocyanate, test for 
iron, 322; thiocyanate, use as test 
for iron, 556. 

Ferrite, defined, a-, B- and j7-, rela- 
tion to tempering, 546; table of 
properties, 546. 

Ferromanganese, use in chilled cast 
iron, 543. 

Ferrous bicarbonate, formation, in 
mineral waters, ores from, 554; 
carbonate, ore, 540; carbonate, 
properties, 554. 

Ferrous chloride, absorption of nitric 
oxide by solutions of, 554; prepara- 
tion, properties, 553; preparation, 
structure, 552. 

Ferrous chromite, 524; preparation 
of potassium chromate from, 527; 
ferricyanide, 321; hydroxide, for- 
mation, properties, 553 ; hydroxide, 
use to reduce indigo, 341; manga- 
nese tungstate, 530; metatantalate, 
523; metacolumbate, 523; oxide, 
553; silicate, formed in metallurgy 
of copper, 429. 

Ferrous sulfate, preparation, proper- 
ties, 554; sulfide, oxidation to 
ferric hydroxide in gas purifiers, 
296; sulfide, preparation, prop- 
erties, use, 556. 

Ferrovanadium, 522. 

Ferrum, 11. 

Fertilizer, calcium cyanamide as, 463 ; 
slag from basic steel process as, 548 ; 
various phosphates in, 461. 

Filters, charcoal, inefficient, 83. 

Fire damp, 286, 287. 

Fireproofing fabrics, with sodium 
tungstate, 530; of cotton goods, 
513; with water glass, 353. 

Fixing in photography, 445. 

Flame colors of calcium, strontium 
and barium, 471; gas and candle, 
structure, 299. 

Flame, oxidizing, blowpipe and Bun- 
sen burner, 304; reducing, blow- 
pipe and Bunsen burner, 304. 

Flames, cause of luminosity, 299; re- 
versed, 304; temperature of, 302; 
source of carbon in}; 299. 

Flint, 348. 

Flint, atomic weight of tellurium, 
190. 

Flint glass, 467. 


ae 


INDEX 


Flour, explosive when mixed with air, 
289 


Flowers of sulfur, 160. 

Fluoboric acid, 367. 

Fluorine, occurrence, 153; 
tion, 153; properties, 154. 

Fluorite, 153, 452. 

Fluosilicic acid, preparation, proper- 
ties, 350. 

Foods, animal, vegetable, inorganic 
constituents, 347; average Ameri- 
can dietary, 347. 

‘“ Fool’s gold,’’ 556. 

Forge, reduction of iron in, 540. 

Formaldehyde, preparation, use as 
disinfectant, 327; use in food for- 
bidden, 328; probable formation 
by reduction of carbon dioxide by 
plants, 312. 

Formalin, see formaldehyde. 

Formic acid, from mercuric fulminate, 
491; preparation from carbon 
monowide, uses, 329. 

Formula of mineral, calculation, 357. 

Formular solutions, definition of, 
183. 

Formulas, meaning of, 16; 
tural, basis for, 323. 

Franklinite, 481. 

Frasch, process for getting sulfur, 161. 

Fraunhofer lines, discovery, 424. 

Free metals, occurrence, 439. . 

Freedom, degrees of, 77. 

Freezing points of solutions and os- 
motic pressure, 360; law for de- 
pression of, 112. 

French Revolution, relation to alkali 
industry, 400; sources of saltpeter 
during, 418. 

Friedel, discovery of carbon in meteo- 
rite from Cafion Diablo, 274. 

Fructose, formation from cane sugar, 
properties, 335; from cane sugar, 
334. 

Fruit, sulfuring of, 164. 

Furnace, electric for aluminium, 495; 
electric for calcium carbide, 462; 
electric for carbon disulfide, 317; 
electric for carborundum, 349; 
electric for phosphorus, 241; re- 
generative, for open hearth steel, 
549, 550. 

Fusible alloys, 269. 


prepara- 


struc- 


Gadolinium, compounds, 505. 

Galena, 161; lead ore, 513; silver in, 
439. 

Gallium, compounds, same as eka- 
aluminium, 506. 

Galvanic cells, calculation of electro- 
motive force, 437; theory, 437. 








581 


Galvanized iron, manufacture, theory 
of conduct, 482. 

Garnet, 349; calculation of formula 
of, 356; orthosilicate, 355. 

Gas, blast furnace; use for heating 
blast, for gas engines, 542; carbon, 
manufacture, uses, 279. 

Gas, correction of volume for changes 
of pressure, 38 ; for temperature, 39. 

Gas, determination of weight of a 
liter of, 40; effect of pressure on, 
34; effect of temperature on, 38; 
effect of water vapor on volume 
and pressure of, 76. 

Gas, illuminating, 295; ‘‘ ideal,’ 94; 
iron pentacarbonyl in, 561; laws, 
graphical representation, 42, 43; 
lighters, cerium-iron, 364; liquors, 
ammoniacal, 202. 

Gas, oil, Pintsch, water, 296; pro- 
ducer, 297; producer, use in re- 
generative furnace for steel, 550. 

Gases, collection and storage of, 22; 
diffusion of, 56; drying of, 54; 
Henry’s law for solution of, 165; 
spectra of, 427; kinetic theory of, 
58; laws of, 34, 38; laws of, exer- 
cises, 43; monatomic, ratio of 
specific heats for, 236; table of 
densities, 95. 

Gasoline, 289. 

Gasometer, description of, 22. 

Gastric juice, hydrochloric acid in, 
406. 

Gayley, dry blast for iron manufac- 
ture, 542. 

Gay Lussac’s law of combining 
volumes, 89; tower, use in manu- 
facture of sulfuric acid, 179. 

Geology, relation of radiochemistry 
to, 475 

Germanium, discovery, compounds, 
361. 

German silver, 560. 

Germicides, sulfites as, 174. 

Gibbs, Willard, phase rule, 107. 

Glass, etching of, 154; manufacture, 
properties, 466; crown, flint, 
strass, paste, Bohemian, Jena, Re- 
sistanz, Non-sol, hard, borosilicate, 
Durax, 467; opaque, stannic oxide 
in, 511; soluble or water, 353. 

Glazes, for earthenware and porce- 
lain, 502. 

Glover tower, use in sulfuric acid 
manufacture, 179. 

Glucinum, name for beryllium, 451. 

Glucose, from hydrolysis of cane 
sugar and of starch, properties, 
334; formation in diabetes, detec- 
tion, 335. 


082 


Gluten from cereals, 8, 335. 
Glycerol, source, use, 326; by-prod- 
uct in manufacture of soap, 322. 

Glyceryl nitrate, 326. 

Gmelin, manufacture of ultramarine, 
502. 

Goiter, connected with deficiency of 
iodine, 144. 

Gold, chlorides of, 450; dioxide, 448 ; 
exercises, 450; monochloride, 450; 
dichloride, 450; trichloride, 45; 
monoxide, 448. 

Gold, occurrence, 445; in sea water, 
washing for, hydraulic mining, 
cyanide process, 446; annual pro- 
duction, coining value, 447; oxides 
of, hydroxide, 448; properties, 
alloys, use of term carat, 448. 

Gold, recovery from copper, 4380; 
trichloride, use in photography, 
445; trioxide, 448. 

Goldschmidt’s thermite process, 497. 

Goldthwaite, jelly making, 337. 

Goodwin, preparation of metallic 
calcium, 452. 

Graebe, discovery of structure of 
alizarin, 341. 

Graham, crystalloids, colloids, 357. 

Graham, study of diffusion of gases, 
57; of liquids, 357. 

Gram, definition, 31. 

Gram atom, definition, 25. 

Gram equivalent, definition of, 184. 

Gram molecular volume, definition, 
94. 

Gram molecule, definition, 25. 

Granites, decomposition to shales, 
clays and soils, 494. 

Graphical representation of gas laws, 
42, 43. 

Graphite, abnormal specific heat, 
397; anodes for alkali manufac- 
ture, 402; heat of combustion, 276 ; 
in cast iron, 543, 544; in colloidal 
solution as lubricant, 277; sources, 
formation, manufacture, proper- 
ties, 276; uses, 276. 

Gravitation, law of, 1. 

Gravity cell, description of, electro- 
motive force of, 437. 

Greenland, cryolite from, 495. 

Green vitriol, 554. 

uo “del Cano, carbon dioxide in, 

Groups of elements, table of, 371. 

Group zero, 236. 

Group I, alternate metals of, general 
properties, 428. 

Group II, alternate metals of, general 
properties, 478; general properties, 
451; solubility of sulfides of, 491. 





INDEX 


Group III, metals of, 494. 
Gaauy IV, metallic elements of, 361, 
62. 

Group V, alternate elements, 522; 
exercises, 272; general properties, 
256; table of oxides, chlorides, 
Se sulfur acids and hydrides, 

ie 

Group VI, alternate elements, 524; 
metallic elements of, 192; proper- 
ties of elements of, 191. 

Group VII, metallic elements of, 156, 
533; missing elements, possible 
reason, 533. 

Group VIII, general statement, 539; 
metals of, general properties, table, 

Guignet’s green, 525. 3 

Gun cotton in smokeless powder, 419; 
manufacture, properties, uses, 338. 

Gun metal, 509. 

Gunpowder, composition, 418; de- 
composition, manufacture, theory 
of burning and explosion, 419. 

Gypsum, 161, 452; plaster of Paris 
from, 457; phase rule, 458; con- 
ditions for formation of, 458, 460; 
vapor pressure of, table, 459. 


Haber, dissociation of hydrobromic 
acid, 146; synthesis of ammonia, 
PAL 

Haddam, columbite from, 523. 

Hafnium, 364. 

Hall, manufacture of aluminium, 495; 
process, 496. 


Halogen acids, table of, 139; family, 


139; meaning of name, 139. 

Halogens, compounds with hydrogen 
and oxygen, 139; general prop- 
erties of, 138; table of, 139. 

Hamburg, cholera in, from water, 
83. 

Hampson, liquid air machine, 234. 

Harcourt, atomic weight of tellu- 
rium, 190. 

Hard glass, 467. 

Hard waters, 310; permanent hard- 
ness, temporary hardness, soften- 
ing by boiling, 463 ; by milk of lime, 
464; by sodium salts, 464. 

Hatchett, discovery of columbium, 
523. 

Haynes, cobalt chromium alloy, 557. 

Heat of combination of chlorine 
with -Hs, .NajivZn; Cu)? Pe 10s- 
of oxygen with Ho, Na, Zn, Cu, P, 
108. 

Heat of combustion, 25; of bitumin- 
ous coal, calculation of, 44; of 
re S, P, Fe, Hg, 27; of hydrogen, 


INDEX 


Heat, mechanical equivalent of, 6; 
units of, 33. 

Helium, coefficient of expansion, 38; 
discovery in sun and in cleveite, 
237; distribution, 238; from 
radium, 9. 

Hematite, 540. 

Hemihedral forms of crystals, 1938. 

Henry’s law, applied to carbon di- 
oxide, 309; for solution of gases, 
165. 

Heptane, 283. 

Heptene, 283. 

Heptine, 283. 

Hexadquochromic chloride, 526. 

Hexagonal System (crystallography), 
194. 

Hexane, 283. 

Hexanitrocellulose, 338. 

Hexathionic acid, 188. 


Hexene, 283. 

Hexine, 283. 

Hillebrand, gases from _ uraninite, 
237; use of sodium pyrosulfate 


for solution of alumina, 408. 

Holmberg, holmium, 505. 

Holmium source, 506. 

Homologue, definition, 289. 

Honey, invert sugar in, 334. 

Hopkins, discovery of Illinium, 505. 

Hornblende, metasilicate, 355. 

Horn silver, 439. 

Hulett, purification of mercury, 485. 

Humidity of rooms, 232. 

Hunyadi water, magnesium sulfate 
in, 480. 

Hydrargyrum, 11. 

Hydrates, defined, 81; of nitric 
acid, 211; of sulfuric acid, 181; 
vapor pressure of, 82. 

Hydration, water of, 82. 

Hydrazine hydrochloride, 222; prep- 


aration, 222: . structure; ©220; 

trinitride, structure, 221. 
Hydriodic acid, constant boiling 

solution, 146; equilibrium with 


hydrogen and iodine, 146; forma- 
tion from potassium iodide, 145; 
heat of formation of, 153; prep- 
aration from iodine, phosphorus 
and water, 145; rate of formation 
and decomposition, 150; reduc- 
tion of sulfuric acid by, 148. 

Hydrobromic acid, constant boiling 
solution of, 143; dissociation of, 
146; preparation from hydrogen 
and bromine, 143; preparation 
from potassium bromide, 142; 
reduction of sulfuric acid by, 142; 
preparation from bromine, phos- 
phorus and water, 143. 








583 


Hydrocarbonate ion, rdle in decom- 
position of carbonates, 375. 

Hydrocarbons, exercises, 305; table 
of, 283. 

Hydrochloric acid, and oxygen, 
equilibrium of reaction between, 
108 ; by-product in making alumin- 
ium oxide from clay, 496; diffi- 
ee with in Leblanc soda process, 

aie 

Hydrochloric acid, constant boiling 
solution of, 120. 

Hydrochloric acid, determination of 
composition by volume, 120; de- 
termination of composition of, 
130; deviation from Boyle’s law, 
35; electrolysis to show composi- 
tion by volume, 120; formation 
from chlorine and hydrogen, 118; 
in gastric juice, 406; oxidation of, 
100. 

Hydrochloric acid, preparation from 
salt and sulfuric acid, 118; prop- 


erties, 119; reactions with hy- 
droxides and oxides, 121; re- 
actions with metals, 120; reac- 


tions with oxidizing agents, 122; 
reaction with potassium perman- 
ganate, 159. 

Hydrocyanic acid, preparation, prop- 
erties, uses, 319. 

Hydrofluoric acid, constant boiling 
solution of, 155; etching of glass 
by, 154; molecular weight, 155; 
preparation, 154; properties, 154. 

Hydrogen, 45; apparatus for prep- 
aration of, 53; burning in chlo- 
rine, 118; chemical properties of, 
59; coefficient of expansion, 38. 

Hydrogen, combination of iodine with 
reversible, 146; combination with 
oxygen reversible, 372; deter- 
mination of atomic weight, 72; 
deviation from Boyle’s law, 35; 
formed in preparing phosphine, 243. 

Hydrogen, formerly unit for atomic 
weights, 68; heat of combustion 
of, 65; ions in liquid ammonia, 
208; in the atmosphere, 45. 

Hydrogen, occurrence, 45; nascent, 
213s palladium semipermeable 
membrane for, 358, 360. 

Hydrogen, preparation, 47; prepara- 
tion by “ hydrone,’’ 52; prepara- 
tion by sodium and potassium, 
50; preparation by ~zine and 
acids, 52; preparation from iron 
and steam, 49; properties of, 55; 
purification of, 54. 

Hydrogen peroxide as_ oxidizing 
agent, 84; as reducing agent, 85; 


584 


preparation, 83; properties, 85; 


uses, 85; structure, 86; tests 
for, 86. 

Hydrogen phosphide, liquid, 243; 
selenide, preparation, properties, 


190; silicide, 349. 

Hydrogen sulfide, decomposition by 
heat, 164; decomposition of solu- 
tion in air, 166; formation from 
elements, 164; ionization of, 168; 
occurrence, 164. 

Hydrogen sulfide, Parsons apparatus 
for generating, 165; reaction with 
iodine, 171; reducing agent, 171; 
removal from hydrogen, 55; re- 
moval from illuminating gas, 295; 
solubility in water, 166. 

Hydrogen telluride, 190. 

Hydrolysis of chlorides, 115; of 
chlorides of phosphorus, 245; 
y salts, theory, 385; of sulfides, 
AA 

‘“ Hydrone,”’ preparation of hydrogen 
by,252. 

Hydronitric acid, preparation, 223; 
properties, 223; structure, 221. 


Hydrosulfate ion, rdle in decomposi- | 


tion of salt, 375. 
Hydrosulfites, see hyposulfites. 
Hydrosulfuric acid, 167. 
Hydrosulfurous acid, 186. 
Hydrotetrachloroantimonic 

267. 

Hydroxide, meaning of name, 21. 
Hydroxides, oxides prepared from, 

392; preparation from metals, 

392; from salts, 393. 
Hydroxylamine, from mercuric ful- 

minate, 491; preparation, 221; 

properties, uses, 222. 
Hydroxylammonium sulfate, 222. 
Hygrometer, moist bulb, 232. 
Hypo-, prefix, 123. 

Hypobromite, sodium, 143. 
Hypochlorites, autoxidation to chlo- 

rates, 125. 

Hypochlorites, preparation, proper- 

ties, uses, 124 
Hypochlorous acid, preparation, 

properties, 124; formation from 

arte and water, 106; structure, 

130. 

Hypochlorous anhydride, 126. 
Hyponitrous acid, preparation, prop- 
erties, 221. 


acid, 


Hypophosphites, use, preparation, 
248; structure, 247. 
Hypophosphoric acid, formation, 
salts, 254. 


Hyposulfite, old name for thiosul- 
fate, 187. 








INDEX 


Hyposulfites, preparation, 186. 
Hypotheses, 2. 


Ianke, George, value of the calorie 
at different temperatures, 33. 

-ic, suffix, meaning, 30; use for acids, 

Ice machines, 204. [123. 

Ice, vapor pressure of, 75. 

-ide, suffix, use, 47. 

Illinium, 505. 

Illuminating gas, ammonia from, 201; 
manufacture, composition, 295; 
whi of hydrogen sulfide from, 
295. 

Imide, definition of, 206. 

Indestructibility of matter, 6. 

India, diamonds from, 275; 
sium nitrate from, 418. 

India rubber as semipermeable mem- 
brane, 358. 

Indicators, acidity or alkalinity of 
at change of color, table, 388; 
chemical nature, 389; choice of an, 
389; definition and list, 122; for 
determining free and combined 
carbonic acid, 464; use of, 387; 
for weak acids and bases, 389. 

Indigo, source, synthesis, use as dye, 
341; use of sodamide in manu- 
facture of, 410. 

Indigo white, 342. 

Indium, discovery, 
compounds, 507. 

Inductive reasoning, 13. 

Infusorial earth, use for dynamite, 
packing and scouring, 351. 

Ingle, separated Bunsen flame, 301. 

Ink, sympathetic, 557. 

Insolubility, effect on a reaction, 376. 

International Bureau of Weights and 
Measures, 31; scale of tempera- 


potas- 


atomic weight, 


Invariant, definition, 78. [tures, 32. 

Invert sugar, 334. 

Iodic acid, 139. 

Iodine and _ starch, 145; com- 
bination with iodine reversible, 
146; in thyroid gland, 144; 


monatomic at high temperatures, 
144; occurrence, 144; liberated by 
nitrous acid, 145; liberated in 
titrating copper, 433; positive in 
iodine trinitride, 224. 

Iodine, properties, 144; reaction with 
hydrogen sulfide, 171; removal 
from hydriodic acid, 145; sodium 
tetrathionate formed by action of 
on thiosulfate, 409; solutions 
standardized by arsenious oxide, 
260; tincture of, 144; trinitride, 
iodine positive in, 223; trinitride, 
preparation, 223. 





INDEX 


fon, definition, 48. 

Ionium, 475. 

Ionization, calculation of degree of 
a freezing points of solutions, 

ihe 

Ionization, degree of, table acids, 
383; bases, 383; salts, 384; effect 
of degree of, neutralization, 384; 
effect on freezing points of solu- 
tions, 112, 381. 

Ionization, evidence of, 112;  ex- 
plained by the electron theory, 182; 
measurement of degree of, 380; of 


acids, 167; of acids, relation to 
hydrolysis of sugar, 381; of 
ammonium hydroxide, 203; of 


compounds of cadmium and mer- 
cury, 491; of first and second 
hydrogen atom of acids, 168; of 
hydrates of chromic chlorides, 526. 

Ionization of oxalic acid, relation to 
solubility of calcium oxalate, 465, 
466; of sulfuric acid, 181; of 
sulfurous acid, 174; of trimethyl 
ammonium hydroxide and tetra- 
methyl ammonium hydroxide, 204 ; 
of water, 171, 383; relation to elec- 
tron theory, 206. 

Iridium, properties, uses, oxides, 
chlorides, double salts, 565; tetra- 
chloride, 565. 

Iron, burning in oxygen, 23. 

Iron carbide, 543, relation to tem- 
pering of steel, 546. 

Iron, dialyzed, 555; discovery «of 
metallurgy of, 390; disulfide, 556; 
heat of combustion, 27; in pro- 
teins, 343; magnetic oxide of 
from burning iron, 23. 

Iron, precipitation of copper by, 435; 
pentacarbonyl, 561; pyrites, 556. 

Iron, reasons for importance of, 539; 
ores, history of use, 540, blast 
furnace, 541; cast iron, 543; 
wrought, 544; cementation and 
cast steel, tempering of steel, 545; 


Bessemer steel, 547; open hearth 
steel, 548; analyses, 544, 551; 
alloy steels, 552; compounds, 552; 
tetracarbonyl, 561. 

Isobutane, structure, 284. 

Isomer, definition, 323; history of 
name, 511. (193. 


Isometric System (crystallography), 
Isothermals of carbon dioxide, 307. 
-ite, use for salts, 124. 

Isotopes, 476. 


Jakowin, reaction of chlorine with 
water, 106. 
James, separation of thulium, 506. 





585 


Jasper, 348. 

Jelly, conditions for making, 337. 
Jena glass, 467. 

Johnson, sherardized iron, 482. 
Joule, defined, 26; relation to calorie, 


Joule-Thomson, effect on expansion 
of compressed gases, 233. 


Kalium, 11. 

Kalk-Stickstoff, 463. 

Kamm, use of potassium silver co- 
baltinitrite in analysis, 559. 

Kaolin, 349; formation, 494; ortho- 
silicate, 355. 

Kekule, formula for benzene, 285. 

ph inventor of process for steel, 
54 

Kelvin, size of molecules, 16; theory 
of the source of oxygen in the air, 
230. 

Kerosene, 
point, 290. 

Ketones, 327. 

Kieselguhr, 351. 


manufacture, flashing 


Kilogram-meter, defined, 32; value 
in ergs, 33. 

Kilowatt, defined, 34. 

Kimberly, diamonds from, 275. 

Kindling temperature, 24. 

Kinetic theory of gases, 58; rela- 


tion to osmotic pressure, 361. 

Kipp generator, 54. 

Kirchoff, discovery of rubidium and 
caesium, 424; discovery of spec- 
trum analysis, 424. 

Knietsch, history of catalytic sulfur 
trioxide, 175. 

Konigsberger, critical temperature of 
mercury, 486. 

Kremann, hydrates of nitric acid, 212. 

Krogh, estimate of coal burned 
annually, 229. 

Krypton, discovery, 238. 

Kurnakow, sodium amalgams, 488. 

Kiister and Kremann, hydrates of 
nitric acid, 212. 


Lacquers, use of nitrocellulose in, 338. 

Lactic acid, formation, structure, 330. 

Lactose, from milk, use, 334. 

Ladenburg, density of ozone, 97. 

Lake Superior region, copper from, 
428. 

Lakes, for dyeing, 501. 

Lampblack, 277. 

Landolt, demonstration of conserva- 
tion of matter, 6. 

Langworthy, respiration calorimeter, 
ois, 

Lanthanum, compounds, 503. 


586. 


Latent heat, former use of term, 74. 

Laudanum, 343. 

Laughing gas, 215. 

Lavoisier, demonstration of the com- 
position of air, 19; determination 
of the composition of air, 227; 
system of nomenclature applied to 
minerals, 356. 

Law, Avogadro’s, 89. 

Law, Boyle’s, 34; Boyle’s, relation to 
Avogadro’s law, 94; Charles, 38; 
Charles, relation to Avogadro’s 
law, 94; Dulong and Petit, 396. 

Law, Faraday’s, 438. 

Law for depression of freezing points 
of solutions, 112. 

Law, natural, 1; of combining vol- 
umes, 89; of combining weights, 
13; of constant proportions, 12; 
of diffusion of gases, 58; of 
gravitation, 1; of ‘‘ mass action,” 
149; of multiple proportions, 87; 
of partial pressures, 41, 77. 

Laws, graphical representation of gas, 
42,43; of gases, 34, 38. 

Lead acetate, preparation, uses, 519; 
basic, 519; carbonate, prepara- 
tion, 519; basic, manufacture, 520; 
comparison with lithopone, 470; 
chloride, formation, solubility, 518; 
chromate, 527; chromate, dis- 
covery of chromium in, 524. 

Lead dioxide, formation, use in 
storage batteries, 516; contrast 
with barium Se 518; glazes, 
danger from, 502; in brass and 
bronze, 431. 

Lead monoxide, manufacture, use, 
515; nitrate, preparation, use, 519. 

Lead, occurrence, metallurgy, 513; 
properties, uses, 514; alloys, 515; 
oxides, 515; compounds, 518. 

Lead oxide, reduction by potassium 
cyanide, 321. 

Lead, treatment by Parke’s process, 
alloys with zinc, 441. 

““ Lead ”’ pencils, 276. 

Lead peroxide, 515; plumbate, red 
lead, preparation, 515; structure, 
516; plumbate, structure, 535; 
red oxide of, 515; sugar of, 519; 
sulfate in storage batteries, 516 ; 
sulfate, preparation, use as pig- 


ment,’ 519; _ sulfide, formation, 
solubility, 518; sulfide, theory of 
precipitation, 169 2 tetrachloride, 


preparation, hydrolysis, 518; tetra- 

sulfate in storage batteries, 518. 
Leather, chrome tanning of, 527. 
eee discovery of soda process, 








‘Lime, 





INDEX 


Leblanc soda process, 411; dis 
covery, 400; recovery of sulfur in, 
456. 


Le Chatelier, explosion waves, 301; 
principle of van’t Hoff-, 111. 

Legislation controling Leblanc soda 
process, 411, 456; manufacture of 
matches, 243; control of lead glazes, 
502; to prevent poisoning by white 
lead, aa le 

Leguminous plants, fixation of nitro- 
gen by, 199. 

Length, unit of, 31. 

Levulose, see fructose. 

Lewis, equilibrium between silver 
oxide, silver and oxygen, 443; 
relation of law of Dulong and Petit 
to Avogadro’s law, 397. 

Liebermann, discovery of structure of 
alizarin, 341. 

Life of an element, 474. 

Lignites, brown and black, composi- 
tion, 280. 
Ligroin, 289. 
Lime kilns, continuous and _inter- 

mittent, 452; slaking of, 453. 

Lime, manufacture, 452; theory of 
formation from calcium carbonate, 
453; in minerals, 357. 

Lime-nitrogen, 463. 

Lime-sulfur wash, 164. 

used to prepare 
alcohol, 325. 

Linde, liquid air machine, 234. 

Liquid air, 232; preparation of 
oxygen from, 20. 

“Liquid smoke, ”’ 
acid, 329. 

Liter, defined, 31. 

Litharge, reduction with blowpipe, 
304; separation of lead from 
silver by formation of, 439. 

Lithium, atomic weight of, 396; 
carbonate, 396; chloride, deter- 
mination of molecular weight of, 
131; comparison with magnesium, 
396; hydride, 395; nitride, 396; 
nitride, formation, 201. 

Lithium, occurrence, properties, 395; 
flame reaction, 396; perchlorate, 
use in determining the atomic 
weight of chlorine, 131; phos- 
phate, 396; urate, 396. 

Lithopone, 470; compared with 
white lead, 521. 

Liversidge, gold in sea water, 446. 

eee discovery of helium in sun, 
VEX 

Louisiana, sulfur in, 161. 

Luminosity of flames, cause of, 299. 

Luna, alchemical name for silver, 444. 


absolute 


contains acetic 


i 


INDEX 


Lunar caustic, 444. 

Lunge, theory of manufacture of sul- 
furic acid by chamber process, 178. 

Lutecium, rare earth metal, 506. 

Luteocobalt chloride, 559. 

Luteorhodium chloride, 563. 


&, definition, 262. 

um, defined, 262. 

Mabery, announcement of Cowles 
furnace, 495. 

McCay, formation of arsenic penta- 
sulfide, 261. 

McCoy, atoms of metallic elements, 
94 


MaclInnes, table for degree of ioni- 
zation, 383, 384. 

Magnalium, 497. 

Magnesia usta, 479. 

Magnesite, 478. 

Magnesium ammonium arsenate, 
259; ammonium chloride, an- 
hydrous magnesium chloride from, 
480; ammonium phosphate, use, 
decomposition to magnesium pyro- 
phosphate, 481; carbonate in hard 
water, 463. 

Magnesium chloride, by-product in 
ammonia soda process, 413; ef- 
fect on the formation of gypsum, 
460; hydrate, conduct on heating, 
480; in salt, 405. 

Magnesium compounds, effect of 
ammonium hydroxide on solutions 
of, 491; diammonium phosphate, 
decomposition, 252; exercises, 493 ; 
hydroxide, formation, theory of 
solubility in solutions of ammo- 
nium salts, 479, 491; metaphos- 
phate, formation, 252.. 

Magnesium nitride, formation, 201; 
occurrence, preparation, proper- 
ties, 478; uses, 479; oxide, prepa- 
ration, uses, 479. 

Magnesium pyrophosphate, forma- 
tion, 252; pyrophosphate, from 
magnesium ammonium phosphate, 
481; silicide, 349; sulfate, 480; 
sulfide, hydrolysis, 491; sulfide, 
preparation, hydrolysis, 480. 

Magnetic oxide of iron, formation 
from iron and steam, 49; from 
burning iron, 23; ore, structure, 
556. 

Magnetite, 540. 

Malachite, 428; formation, 431. 

as acid, carbon suboxide from, 

16. 

Maltodextrin, from starch, 334. 

Maltose, alcohol from, 325; forma- 
tion from starch, hydrolysis, 334. 








587 


Mammoth cave, saltpeter from, 418. 

Manchot, combination of ferrous sul- 
fate with nitric oxide, 217; ferrous 
chloride and nitric oxide, 554. 

Manganates, preparation, 536. 

Manganese dioxide, history of uses, 
535; use in preparing oxygen, 21. 

Manganese heptoxide, properties, 
538; list of oxides, 538; occur- 
rence, properties, 533; alloys, uses, 
compounds, 534; valence, structure 
of compounds, 534; tetrachloride, 
probable formation, 101. 

Manganic acid, change to perman- 
ganic acid, 536. 

Manganous chloride, 530. 

Manganous hydroxide, 535. 

Manganous manganic oxide, 534, 
structure, 535. ? 

Manganous sulfides, 535. 

Manometer, 41. 

Maple sugar, 333. 

Marsh’s test for arsenic, 257. 

‘“ Mass action,” law of, 149. 

Mass and weight, relation, 32; de- 
pendent on velocity, 5. 

Matches, 242; from tetraphosphorus 
trisulfide, 243; law forbidding 
ordinary phosphorus in, 243. 

Matte, copper, 429; nickel, 559. 

Matter, conservation of, 6; defini- 
tion, 5; indestructibility of, 6. 

Mauve, discovery of, 340. 

Mechanical energy, units of, 
equivalent of heat, 6. 

Medicine, relation of radiochemistry 
to, 475. 

Meerschaum, trisilicate, 356. 

Meker burner, temperature of, 303. 

Melting point, criterion of pure sub- 
stance, 12. 

Melting points of elements, 372, 


Sol; 


table, 373; in absolute tempera- 
ture, 135. 
Mendeleef, identification of scan- 


dium, as ekaboron, 503; gallium as 
eka-aluminium, 506; periodic sys- 
tem, 136. 

Mental work in respiration calorim- 
eter, 316. 

Menzies, critical temperature of mer- 


cury, 485. 

Mercuric chloride, ionization of 
anomalous, 382; chloride, prep- 
aration, 489; properties, uses, 
antidote for, 490; cyanide, prep- 
aration, decomposition, 490; de- 


rivatives of ammonia, 492. 
Mercuric fulminate, use, hydrolysis 

to hydroxylamine, 491; iodide, for- 

mation, complex compound with 


588 


potassium iodide, use, 490; iodide, 
Nessler’s reagent from, 492; ni- 
trate, 490. 

Mercuric oxide, decomposition, 9; 
formation and decomposition, 19; 
preparation, 488; yellow, 489. 

Mercuric sulfide, red and black, use, 
489; solubility, 491. 

Mercurous chloride, formation in 
amalgamation process, 441; chlo- 
ride, preparation, vapor density, 
formula, uses, 489; nitrate, prep- 
aration, oxidation and reduction 
of, 490; basic, 490; oxide, 488; 
sulfate, use in Weston and Clark 
cells, 437. 

Mercury, abnormal ionization of 
compounds of, 491; exercises, 493 ; 
*heat of combustion, 27; occur- 
rence, 484; metallurgy, purifica- 
tion, properties, critical tempera- 
ture, 485; uses, amalgams, 486; 
use in recovering gold, 446; use 
in Castner-Kellner process, 402. 

Metaboric acid, 366. 

Metachloroantimonates, 268. 

Metallic elements in periodic system, 
136. 

Metallurgy, development of, 390; 
electrolytic methods in, 391; of 
aluminium, history, 391; roasting 
of sulfides in, 391; use of fuels in, 


390. , 
Metals, characteristics, 369. 
Metals, classification, 370, 371; 


preparation by thermite process, 


495; spectra of, 427; systematic 
study of, 390. 
Metantimonic acid, 267. 
Metaphosphoric acid, formation 


properties, 253; hydrolysis, 249; 
properties, polymeric forms, 254. 

Metasilicates, 355. 

Metasilicic acid, 355. 

Metastannic acid, preparation, prop- 
erties, 512. 

Metastannyl chloride, 512. 

Metathesis, defined, 81. 

Meteorites, iron and nickel in, 540. 

Meter, definition, 31. 

Methane, kindling temperature, 288. 

Methane, limits for explosive mix- 
ture with air, 288; occurrence, 
preparation from sodium acetate, 
properties, 286; structure, 284; 
substitution products of, 287. 

Methyl alcohol, manufacture, 324, 
properties, use, 325; relation to 
structure of methyl ether, 324. 

Methyl amine, 339. 

Methyl ammonium iodide, 339. 








INDEX 


Methyl chloride, from methane, 287. 

pans chloride, from methane, 
287. 

Methyl ether, determination of struc- 
ture, 324. 

Methyl iodide, relation to structure of 
methyl ether, 324. . 

Methyl red, use in determining free 
carbonic acid, 465. 

Methyl silicate, colloidal silicic acid 
from, 353. 

Meyer, Lothar, preparation of hy- 
driodic acid, 145. 

Meyer, V., kindling temperature of 
methane, 288. 

Mho, defined, 380. 

Mica, 348; orthosilicate, 355. 

Michael, explosion waves, 301. 

Michelson, explosion waves, 301. 

Michigan, bromine from brines in, 
140. 

Microcosmic salt, use, 423; formula, 
decomposition, use, 253. 

Micron, defined, 262. 

Migration of ions in electrolysis, 113. 

Milk sugar, source, use, 334. 

Millikan, number of molecules in I1cc., 
16, 96. 

Milliliter, used instead of cc., 31. 

Mineral, calculation of formula of, 
356. 

Mirrors, tin amalgam for, 487. 

Mispickel, 256. 

Mixer for iron, 542. 

Mixtures and pure substances, 7. 

Moissan, artificial diamonds, 274; 
preparation of fluorine, 153. 

Moisture, determination of in air 
by weighing, and dew point, 2382; 
effect on chemical reactions, 312; 
precipitation from air at an alti- 
tude, 232; presence in air, 231 

Mol, definition, 183. 

Molar solutions, definition of, 183. 
Molecular weights, determined by 
measuring osmotic pressure, 360. 
Molecules of the elements, 93; 

number of in 1 cc., 95, 16. 

Molybdenite, 528. 

Molybdenum, occurrence, properties, 
528; compounds, 529; _ trioxide, 
properties, 528; use in molybdic 
solution, 529. 

Molybdic anhydride, complex com- 
pounds from, 529. 

‘‘Molybdic solution’’ preparation, 
use to determine phosphoric acid, 
529, 530. 

Molybdic sulfate, reduction of molyb- 
dic acid to for phosphorus de- 
terminations, 529. 


OO 


INDEX 


Monatomic gases, ratio of specific 
heats, 236. 

Monazite sand, cerium from, 363, 
thorium from, 364. 

Monocalcium phosphate, 249; 
bility, hydrolysis, 461. 

ar ae system (crystallography), 
195. 

Monopotassium diarsenite, 259. 

Monosodium phosphate, 249; methyl 
orange as indicator for, 251. 

Montana, arsenic from smelting fur- 
nace in, 256. 

Moore, dissociation and ionization 
of ammonium hydroxide, 204. 

Mordants, 342; potassium pyrochro- 
mate, 527; titanium compounds as, 
363. 

Morley, determination of the compo- 
sition of water by volume, 68; by 
See oxygen and hydrogen, 

il 


Morphine, 343. 

Mortar, composition, hardening of, 
454. 

Mother of vinegar, 329. 

Muffle furnace, 440. 

Multiple-effect evaporators, 405. 

Multiple proportion, law of, 87. 

Munroe-Neubauer crucibles, 565. 

Muscular energy in respiration calo- 
rimeter, 315; 346. 

Musgrave, development of Leblanc 
soda process, 400. 


Naphtha, 289. 

Naphthalene, from coal tar, use, 295. 
‘* Nascent,”’ definition, 213. 
Natrium, 11. 

Natural gas, 286. 

Natural law, definition, 1. 


solu- 


Nature of chemical energy, 27; of 
scientific knowledge, 1. 

Negatives in photography, 445. 

Neodymium, discovery, separation 


from praseodymium, compounds, 
504. 

Neon, discovery, 238. 

Nernst, equilibrium for formation of 
nitric oxide, 216; quantum theory, 
398 ; table for periodic system, 135; 
lamp, 363. 

Nessler’s reagent, 490, 492. 

Neubauer-Munroe crucibles, 565. 

Neutrality, definition of, 385. 

Neutralization, definition, 121; in 
liquid ammonia solutions by union 
of hydrogen and amide ions, 208; 
theory of imperfect, 386, 384. 

New Caledonia, nickel from, 559. 

Newton’s law of inertia, probable 











589 


basis for principle of van’t Hoff-Le 
Chatelier, 111. 

Nickel carbonyl, preparation, proper- 
ties, 561; chloride, 560; dimethyl- 
glyoxime, separation from cobalt, 
560; occurrence, properties, 559; 
uses, alloys, compounds, 560; sul- 
fate, 560. 

Nicotine, 342. 

Niobium, see Columbium. 

Niton, discovery, 238; 
A474; half-life of, 474. 

Nitrates, formation in soils, 
oxides prepared from, 392. 

Nitric acid, action on copper, 213; 
action on zine, 213; and nitric 
oxide from nitrogen peroxide, 220; 
as oxidizing agent, 212; decom- 
position of, 212; detection with 
ferrous sulfate, 217; formation of 
nitrates from, 212. 

Nitric acid, hydrates of, 211; oxida- 
tion of feathers or wool by, 212; 
preparation from sodium nitrate, 
210; properties, 212; structure 
according to electron theory, 207. 

Nitric oxide and carbon bisulfide, 
flame of, 217; and ferrous chloride, 
554; and nitric acid from nitro- 
gen dioxide, 220; coefficient of ex- 
pansion, 38; combination with 
ferrous sulfate, 217; deviation 
from Boyle’s law, 35; equilibrium 
of formation, 215; formation of 
nitrites from, 217; formed by re- 
duction of nitric acid with ferrous 
sulfate, 217; formed by union of 
nitrogen and oxygen, 215; prepa- 
ration by action of nitric acid on 
copper, 215; structure, 217; use 
in manufacture of sulfuric acid, 
L733 

Nitrobenzene, use in preparing ani- 
line, 340. 

Nitrogen, coefficient of expansion, 38 ; 


properties, 


199: 


combination with oxygen, 200; 
with hydrogen, 201; deviation 
from Boyle’s law, 35; exercises, 


225; iodide, preparation, proper- 
ties, 225; liquid, 234; list of 
oxides of, 214; pentoxide, prepa- 
ration, properties, 220; dioxide, 
formation from nitric oxide, 219; 
dioxide, from copper nitrate, 434; 
dioxide, from nitric oxide, 217; 
ratio of specific heats of, 237; tri- 
chloride, endothermic, 225;  tri- 
chloride, equivalent to 3 Cle, 224; 
trichloride, formation by action of 
chlorine on ammonia, 209; occur- 
rence and natural history, 198; 


590 


trichloride, oxidation of arsenious 
oxide by, 224; trichloride, prepara- 
tion, properties, 224; peroxide, 
structure, 219; peroxide, use in 
manufacture of sulfuric acid, 178; 
preparation from air, 200; prepara- 
tion from ammonium nitrite, 200; 
from sodium nitrite, 200; proper- 
ties, 200; sources of for vegetable 
growth, 198; tetroxide, dissocia- 
tion, 219; tetroxide, formation of 
nitrous and nitric acids from, 219; 
tetroxide, relation to nitrogen per- 
oxide, 219; tetroxide, structure, 219. 

Nitroglycerin, 326; explosion of, 327. 

Nitro Nitrogen Trichloride, probable 
existence, 225; hydrolysis of, 226. 

Nitrosyl chloride, properties and 
hydrolysis, 214. 

Nitrosylsulfuric acid, formation in 
manufacture of sulfuric acid, 178. 

Nitrous acid, formation, 218; anhy- 
dride, dissociation, 218; anhydride, 
preparation, 218; anhydride, use in 
manufacture of sulfuric acid, 178; 
oxide, deviation from Boyle’s law, 
35; oxide, preparation, properties, 
structure, uses, 214; oxide, use as 
anesthetic, 215. 

Nomenclature, lLavoisier’s system 
applied to minerals, 356; of acids 
and salts, 123; of binary com- 
pounds, 29; of oxides, 29. 

Noncoking coals, 281. 

Nonmetallic elements 
system, 136. 

Non-metals, characteristics, 369. 

Non-sol glass, 467. 

Nordhausen sulfuric acid, 556. 

Normal solutions, definition of, 184. 

porten. composition of silicic acids, 

Ne ay manufacture of nitrates in, 

Noyes, A. A., cause of slight solu- 
eae of cobalt and nickel sulfides, 

Noyes, W. A., determination of the 
composition of water, 71. 

Number of molecules in 1 cc., 16. 

Nutrition, diet, 345, 347; studied 
with respiration calorimeter, 345. 


in periodic 


Octahedron (crystal), 193. 

Octane, 283. 

Octene, 283. 

Octine, 283. 

Octovalent, definition, 64. 

Ohm defined, 33. 

Oil gas, 296; percentage composi- 
tion, 299. 





INDEX 


Oil of vitriol, 46. 

Oleic acid, source, 331. 

Olein, 332. 

Opal, composition, 354. 

Open hearth steel, 548; regenerative 
furnaces for, 549, 550. 

Opium, 343. 

Ordinates, axis of, 43. 
Se assay of for gold and silver, 
40. 
Origin, mathematical definition, 43. 
Orpiment, 256; preparation, proper- 
ties, uses, 260. 

Orthoantimonic acid, 267. 

Orthoclase, trisilicate, 356. 

Orthophosphoric acid, classes of salts, 
249; decomposition of salts of, 
252; formation, 248; ionization, 
250; preparation, 249; properties, 
249; solubility of salts of, 252; 
structure, 247. 

Orthosilicates, 355. 

Orthosilicic acid, 355. 

Oscillators, in quantum theory, 398. 

““ Osmic acid,’’ see osmium tetroxide, 
565. 

Osmium, catalyzer for synthesis of 
ammonia, 201. 

Osmium-iridium, composition, prop- 
erties, 564; ruthenium in, 563. 
Osmium, occurrence, oxides, 564; 

chlorides, osmates, 565; tetroxide, 
us? in histology, 565. 
Osmosis, 358. 


Osmotic pressure, connection with 


freezing points and boiling points 
of solutions, 360; defined, 360; 
measurement of, 359. 

-ous, suffix, meaning, 30; 
acids, 123. 

Oxalic acid and hydrazine from bis- 
diazoacetic acid, 222; carbon mo- 
noxide from, 311; decomposition, 
466; ionization, relation to solu- 
bility of calcium oxalate, 465, 466; 
occurrence, manufacture from so- 
dium formate, 329; strength illus- 
Bae 386; use, decomposition, 

Oxidation, definition, 63; reactions 
for potassium permanganate, 537; 
no eae equations for reactions of, 
Wel 

Oxide of mercury, decomposition, 9. 

Oxides, nomenclature of, 29; of 
nitrogen, summary of methods of 
preparation, 214; preparation from 
metals, nitrates, carbonates and 
hydroxides, 392; by _ precipita- 
poe 393; valence of elements in, 
157. 


use for 


INDEX 


Oxidizing flame, blowpipe and Bun- 
sen burner, 304. 

Oximes, 222. 

“Oxone,”’ preparation of oxygen 


from, 21. 

Oxyacids of chlorine, structure, 
130. 

Oxygen, 19; absorbed by molten 
silver, 448; and acid properties, 
23; and chlorine, comparison of 


heats of combination, 108; basis 

of unit for atomic weights, 68; 

coefficient of expansion, 38; com- 

pees with hydrogen reversible, 
2. 

Oxygen, determination of in air, 227; 
deviation from Boyle’s law, 35; 
dissociation pressure of from 
barium peroxide and manufacture, 
469; for medicinal use from liquid 
air, 235; liquid, 234; occurrence, 
19; of air, may have come from 
carbon dioxide, 230; origin of 
name, 23. 

Oxygen, preparation from liquid air, 
20; mercuric oxide, 19; ‘‘ oxone,”’ 
21; potassium chlorate, 20; potas- 
sium chlorate with manganese 
dioxide, 21; sodium peroxide, 21. 

Oxygen, properties of, 22; weight 
of one liter in different latitudes, 
Dade 

Oxyhydrogen blowpipe, 61. 

Ozone, action on silver, 443; prepa- 
ration, 97; properties, 97; struc- 


ture, 98; from action of fluorine 
on water, 154. 

Palladium, semipermeable mem- 
brane for hydrogen, 358, 360; 


catalytic effect, use as catalyzer 
for reduction of fats, 564; di- 
chloride, use in gas analysis, 564; 
occurrence, properties, absorption 
of hydrogen, 563; oxides, chlo- 
rides, ammines, 564. 

Palmaer, absolute potential of ele- 
ments, 436. 

Palmitic acid, source, 331. 

Palmitin, 331. 

Paper, manufacture, 337; 
338. 

Paraffin, 290. 

Parastannic acid, 512. 

Paregoric, 343. 

Paris green, 259. 

Parke’s process for silver, 440. 

Parsons’ apparatus for hydrogen 
sulfide, 165. 

Partial pressures, law of, 41, 77. 

Paste, glass, 467. 


sizing, 





591 


Pattison’s process for silver, 439, 

Peat, composition, 280. 

Pectin, relation to jelly, 337. 

Pectose, relation to jelly, 337. 

Pens, iridium for tips of, 565. 

Pentane, 283. 

Pentathionic acid, 188. 

Pentine, 283. 

Pepsin, 344. 

Per-, prefix, meaning, 30, 123. 

Perchlorates, 128. 

Perchloric acid, 128; structure, 
130; structure of hydrated, 129. 

Perchloric anhydride, 129. 

Periodic law, exceptions to, 138. 

Periodic system, 132; tables, 134, 
135, 

Perkin, discovery of mauve, 340. 

Perkin, fireproofing of cotton goods, 
513. 

Permanent hardness, 311. 

Permanganates, preparation, 537. 

Permanganic acid, formed by use of 
sodium bismuthate, 269; from 
manganic acid, 537. 

Permanganic anhydride, 538. 

Permonosulfuric acid, 188. 

Peroxides, structure, 518. 

Perrin, estimate of number of mole- 
cules in 1 cc., 96. 

Persulfuric acid, preparation, uses, 
187. 

Pertitanic acid, 
titanium, 362. 

Petrolatum, 290. 

Petroleum ether, 289. 


use in detecting 


Petroleum, occurrence, localities, 
varieties, refining, 289. 
Pewter, 509. 


Pharmaceutical extracts, use of al- 
cohol in, 325. 

Phase, effect of escape of a gaseous, 
376; solid, effect of on reaction, 
Bits 

Phases, definition, 77; 
hydrate, 107. 

Phase rule, 107; dissociation of 
calcium carbonate, 453; plaster of 
paris, 458 ; transition or quadruple 
point for sodium sulfate, 406. 

Phenacetine, 340. 

Phenol, source, manufacture, prop- 
erties, use as antiseptic, 326. 

Phenolphthalein, use as indicator in 
liquid ammonia, 208; use in 
determining free carbonic acid, 
464. 

Phenylhydrazine, derivative of hydra. 
zine, 223. 

Phosgene, 316. 

Phosphate rock, 241. 


of chlorine 


592 


Phosphine, compared with ammonia, 
arsine and stibine, 244; prepara- 
tion, properties, 243. 

Phosphomolybdic acids, 529. 

Phosphonium group, unstable, 244; 
iodide, interference in preparing 
hydriodic acid, 145; iodide, prep- 
aration, 243. 

Phosphor bronze, 431. 

Phosphoric acid, determination with 
molybdic solution, 529, 580; 
formed by burning phosphine, 
243; water soluble, citrate-solu- 
ble, insoluble, in fertilizers, 461. 

Phosphorous acid, preparation, prop- 
erties, 248; structure, 247. 

Phosphorus, acids of, list, 247; 
acids of, basicity, 247; allotropic 
forms, 241; burning in oxygen, 
23; chlorides of, hydrolysis, 245; 
exercises, 255. 

Phosphorus, heat of combustion, 27; 
in proteins, 343; occurrence, 240; 
oxides of, 246; oxychloride, 
formed by hydrolysis of the penta- 
chloride, 245; oxychloride, prep- 
aration from the trichloride, 246. 

Phosphorus pentachloride, action ou 
hydroxyl compounds, 245; dis- 
sociation, 245; hydrolysis, 116, 
245; preparation, properties, 244. 

Phosphorus pentasulfide, properties, 
uses, 254. 

Phosphorus pentoxide, efficiency as 
drying agent, 246; from burning 
phosphorus, 23; moisture left 
in gas by, 54; preparation, prop- 
erties, 246. 

Phosphorus, positive and negative 
valences, 248; preparation, 241; 
red, 241; sulfides of, 254; sulfide, 
use for matches, 243. 

Phosphorus tetroxide, 246;  tet- 
roxide, hydrolysis, 254;  trichlo- 
ride, hydrolysis, 115; trichloride, 
hydrolysis, 245; trichloride, prep- 
aration, properties, 244 “triox- 
ide, preparation, properties, 246; 
valence in acids, 248; yellow, 241. 

Phosphotungstic acid, 531. 

Photographic plate, effect of Réntgen 
and Becquerel rays on, 471. 

Photography, 444, dry plates, posi- 
tives, negatives, developing, fixing, 
toning,, 445. 

Photometry, stellar, use of selenium 
in, 190. 

Photosphere of sun, spectrum, 426. 

Physical sciences, 4. 

Physics, definition, 5. 

Pictet, liquefaction of air, 233. 


INDEX 


Pig iron, continuous casting ma- 
chines for, composition, gray, 
white, chilled, 543; analyses, 544. 

Pintsch gas, 296. 

Pitchblende, uranium in, 531. 

Plank, quantum theory, 398. 

Plaster of Paris in cement, 454 ; 
manufacture, use, 457; phase 
rule, 458. 

Plating, silver, 442. 

Platinic chloride, preparation, prop- 
erties, 566. 


Platinized asbestos, catalysis of 
union of O and H by, 62; prep- 
aration, 62. 

Platinous chloride, 565. 

Platinum ammines, 566; catalyzer 
for sulfur dioxide, 175; disulfide, 


566; iridium electrodes in alkali 
manufacture, 401; metals, gen- 
eral properties, table, 562; prints 


in photography, 445; properties, 
uses, catalytic action, sponge, 
565. 


Pliicker tubes for spectra of gases, 
ADT. 

Plumbic acid, 516. 

Plumbum, 11. 

Poisoning by white lead, 521; by 
lead water pipes, 514. 

Polarimeter, used to determine sugar, 
3o0% 

Polarization of light by crystals, 196. 

Polarized light, effect of sugar on, 
333. 

Polysulfides, ammonium, 422. 

Polythionic acids, 188. 

Porcelain, 501; glazing, 502. 

Positives in photography, 445. 

Potassium aluminium sulfate, 500; 
argenticyanide, 320; use in silver 
plating, 321; aurate, 450; aurous 
cyanide, formation in cyanide 
process, 446; bicarbonate, prep- 
aration, use, 420. 

Potassium carbonate from wood 
ashes, 414; from beet sugar manu- 
facture, from wool, 419; prop- 
erties, 419; uses, 420. 

Potassium chlorate, composition, 21; 
manufactuie, uses, 416; prepa- 
ration, 127; preparation of oxygen 
from, 20. 

Potassium chloroaurate, 450. 

Potassium chloride, measurement 
of degree of ionization of, 380; 
conductance of solutions of, 380; 
‘properties, use in fertilizers, 416; 
for manufacture of saltpeter, 416, 
418; chloroplatinate, use in de- 
termining atomic weight of chlo- 


Se 


INDEX 


rine, 130; chloroplatinate, 566; 
chloroplatinite, use in photogra- 
phy, 565; chloroplumbate, 519; 
chromate, ‘preparation from chrome 
iron ore, 527; chromium sulfate, 
527 ; cobaltinitrite, formation, prop- 
erties, use in analysis, 558, 419; 
colbalticyanide, 558; cobaltocyan- 
ide, 558; cupric chloride, use in 
iron analysis, 432; cuprocyanide, 
435; cyanate, 321; cyanide, for- 
mation, preparation, use, 319 ; cyan- 
ide, preparation, use, 420. 

Potassium diuranate, 532; dichro- 
mate, mordant, 342; dichromate, 
reduction by hydrogen = sulfide, 
171; ferrate, 553; ferricyanide, 
formation from potassium cya- 
nide, 320; ferricyanide, use in 
blue-print paper, 331; ferric ferro- 
cyanide, 321; ferrocyanide, forma- 
tion from potassium cyanide, 320; 
Prussian blue from, 320 ; ferro- 
cyanide, preparation, 319; flu- 
oride, acid, 155; fluotantalate, 
use in purifying tantalum, 524. 

Potassiumhydroxide, preparationfrom 
the carbonate, 415; by electroly- 
sis, properties, 415; slight effect 
on glass, use in analysis, 416. 

Potassium iodate, reduction in pre- 
paring the iodide, 417; iodide, 
manufacture, uses, 417. 

Potassium manganate, preparation, 
conversion to permanganate, 536; 
mercuric iodide, 490; metachlo- 
roantimonate, 268; metallic, dis- 
covery, preparation, properties, 415. 

Potassium nitrate, formation in 
soil, 199; nitrate, relation to 
gunpowder, 417; sources, manu- 
facture, properties, uses, 418; 
nitrite, preparation, properties, 
use, 419. 

Potassium, occurrence, relation to 
minerals, clays, soils, plant growth, 
414; osmate, 565; oxide, 415; 
perchlorate, preparation, 128; per- 
chlorate, preparation, properties, 
416. 

Potassium permanganate, properties, 
uses, 537; typical oxidations with 
537; reaction of hydrochloric 
acid with, 159; purification of 
hydrogen by, 55. 

Potassium perruthenate, 563; per- 
sulfate, preparation, 187;  polyi- 
odides, formation, use in iodimetry, 
417; pyrochromate, preparation, 
properties, uses, relation to pyro- 
sulfate, 527; use in chrome tanning, 





593 


527; pyrosulfate, decomposition, 
use in analysis, 417. 


Potassium, retention in soils by colloi- 


dal silicic acids, 354; ruthenate, 
563; silver cobaltinitrite, 559, 
419: silver cyanide, use in silver 


plating, 442. 

Potassium, source for shales, clays 
and soils, 494; sulfate, 417; 
sulfate, acid, preparation, pyro- 
sulfate from, 417; sulfocarbonate, 
preparation, 317, use, 318; tar- 
trate, acid, 330; tetroxalate, use 
as standard in alkalimetry, 330. 

Potassium thiocyanate, carbon oxy- 


sulfide from, 318; thiocyanate, 
preparation, use in testing for 
iron, 321; triiodide, 145; zincate, 


483. 

Potential, absolute, of elements, 436; 
differences of in relation to corro- 
sion of iron, 550. 

Power, unit of, 33. 

Praseodymium, discovery, separa- 
aoe from neodymium, compounds, 
504. 

Precipitation, theory of, 376. 

Prenner and Schupp, molecular weight 
of sulfur vapor, 163. 

Preparation of compounds, general 
methods, 372-379; of pure sub- 
stances, 8. 

Pressure, effect of on a gas, 34. 

Priestly, analysis of air by nitric 
oxide, 230. 

Primary salts, 249. 

Printing in photography, 445. 

Producer gas, 297; heat relations 
in manufacture, 598 ; percentage 
composition, 299 ; use in regenera- 
tive furnace for steel, 550. 

Propane, structure, 284. 

Propene, 283; structure, 284. 

Propine, 283. 

Propylene, structure, 284. 


Proteins, occurrence, 343; diges- 
tion, 344. 

Prussian blue, 320; use in chrome 
green, 525. 


Prussic acid, 319. 

Ptomaines, 343. 

Ptyalin, 344. 

Pure substances and mixtures, 7; 
composition of expressed in mul- 


tiples of atomic weights, 17; 
distinguished from mixtures, 12; 
preparation of, 8. 

Purpureocobalt chloride, 559. 

Pyridine solutions with semiper- 
meable membrane, 358. 

Pyrite, 161. 


594 


Pyrite burners, 179; use of oxide of 
iron from, 540. 

Pyroantimonic acid, 267. 

Pyroarsenic acid, 259. 

Pyroboric acid, 367. 

Pyrolusite, 533. 

Pyrophosphoric acid, hydrolysis, 248; 
preparation, properties, 253. 

Pyrosulfates, preparation, 186. 

Pyrosulfuric acid, 186. 


Quadrivalent, definition, 64. 
Quadruple point for sodium sulfate, 
406 


Qualitative analysis, basis for groups 
of, 166; definition, 66; groups of, 
166; of water, 66. 

Qualitative synthesis of water, 66. 

Quantitative analysis, definition, 66. 

Quantitative synthesis of water by 
volume, 66. 

Quantum theory, 398. 

Quartz, 348; properties, fused, uses, 
352. 

Quinine, 343. 

Quinquivalent, definition, 64. 


Radiations, penetrating, 471; kinds 
of, 472, 473. 

Radical, definition, 47; definition, 
relation to structure, 323. 

Radioactive elements, series of, 475. 

Radiochemistry, relation to geology 
and medicine, 475. 

Radiothorium, 364. 

Radium, an element, 9, 473; chemical 
action, 475; discovery, 471; dis- 
integration, 473; emanation, 474; 
evolution of heat by, 472; half- 
life, 474; nature of rays, 473; 
properties, 472. 

Rails, steel, manufacture, 548. 

Ramsay, discovery of argon, 235; 
discovery of helium, 237 ; discovery 
of helium from radium, 473; disso- 
ciation of nitrous anhydride, 218; 
possible disintegration of atoms by 
radium emanation, properties of 
niton, 474, 475; use of periodic sys- 
tem in discovery of noble gases, 136. 

Rare earth elements, position in 
periodic system, 134, 138. 

Rare earths, general, 502; groups of, 
rape methods of separation, 503, 

Raschig,' theory of 
manufacture, 179. 

Rayleigh, discovery of argon, 235. 

Rays, 4, 8, 473; Y, 6, 474. 

Bayes chemical action of radioactive, 


sulfuric acid 





INDEX 


Reacting substances, effect of remov- 
ing one on equilibrium, 152. 

Reactions, bimolecular and unimolec- 
ular, 150; calculation of relative 


speed of at equilibrium, 151; 
effect of insolubility on, 376; 
effect of volatility on, 374; re- 


versible, 50; 
Pie 

Realgar, 256; preparation, proper- 
ties, uses, 260 

Reasoning, inductive, 13. 

Reciprocal ohms, definition, 380. 

Red lead, oxide, manufacture, 515; 
structure, 516. 

Reducing agent, sodium amalgam, 
zinc amalgam, 487 ; flame, blowpipe 
and Bunsen burner, 304. 

Reduction, definition, 63; writing 
equations for reactions of, 171. 

Refining, electrolytic, of copper, 429. 

Refrigeration by ammonia in ma- 
chines, 204. 

Regenerative furnace for open hearth 
steel, 549, 550. 

Regular system (crystallography), 
193. 

Reid, discovery of indium, 506. 

Resistanz glass, 467. 

Resonators in quantum theory, 398. 

Respiration calorimeter, 313; study 
of nutrition with, 345. 

Reversed flames, 304. 

Reversible reactions, 50; hydro- 
chloric acid and oxygen, 109; 
hydrogen and iodine, 146; ioniza- 


speed of chemical, 


tion, 115; salt and sulfuric acid, 
119; theoretically all reactions, 
Sin 
Rhodium, properties, oxides, chlo- 
rides, complex salts, 563. 
Rhombic dodecahedron (crystal), 
193; hexahedron (crystal), 195; 


system (crystallography), 194. 
Richter, discovery of indium, 506. 
Roasting, defined, 263; sulfides in 

metallurgy, 391. 

Rochelle salt, use in Fehling’s solu- 

tion, 335. 

Rock crystal, use for lenses, 352. 
Rock salt, mining and obtaining, 404. 
See discovery of Réntgen rays, 

A471. 

Rosa, respiration calorimeter, 313. 
Roscoe and Schorlemmer, Treatise, 

standard for ventilation, 231. 
Rose, distinguished columbium and 

tantalum, 523. 

Roseocobalt chloride, 559. 
Roseorhodium chloride, 563. 
Rouge, 555. 


INDEX 


Rubidium alum, 424; chloroplatinate 
424, 393; discovery, occurrence 
in carnallite, properties, 424. 

Ruby, 494; artificial, 500. 

Rum, 325. 

Ruthenium, occurrence, oxides, chlo- 
rides, double chlorides, ruthenates, 
perruthenates, 563. 

Rutherford, number of molecules in 
oe 16, 96; disintegration theory, 

2: 


Saccharimeters, 333. 

Sackur, quantum theory, 398. 

Safety lamp, Davy. 

Safety matches, 242. 

Sainte-Claire-Deville, preparation of 
aluminium, 495. 

Saleratus, 420. 

Sal soda, 411. 

Salt and sulfuric acid, reversible re- 
action, 374; definition, 47, 121; 
mining and obtaining, 404. 

Saltpeter, Chili, 210; formation in 
soil, 199. 

Salts, nomenclature of, 124; hy- 
drolysis, general statement, 394; 
list of insoluble, 393; solubility of, 


general statement, 393; valence 
of elements in, 158. 
Samarium, discovery, compounds, 


505. 

Sapphire, 494, 500. 

Scandium, discovery, same as eka- 
boron, compounds, 503; same as 
-ekaboron, 136. 

Scheele, discovery of chlorine, 535. 

Science, subdivisions of, 3. 

Sciences, abstract, physical, bio- 
logical, psychological, 4. 

Scientific knowledge, Nature of, 1. 

Scott, determination of the composi- 
tion of water by volume, 68. 

Sea water, gold in, 446. 

Secondary salts, 249. 

Selection of atomic weights, 16, 92. 

aoc acid, preparation, properties, 
190. 

Selenite, 457. 

Selenium dioxide, decomposition by 


sulfur dioxide, 190; preparation, 
190. 
Selenium, occurrence, allotropic 


forms, uses, 189. 
Self-hardening steels, 552. 
noe coals, composition, 
Semipermeable membranes, 357; 
preparation, 358; mechanism of 
action, 358, 360; for pyridine, 
water, hydrogen, 358. 





595 


Septivalent, definition, 64. 
Series of hydrocarbons, 283; 
tromotive, 435, table, 
radioactive elements, 475. 
Serpentine, disilicate; 356. 
Sexivalent, definition, 64. 
Shales, formation of, 494. 
Sherardized iron, 482. 
Sicily, sulfur in, 160. 
Siderite, 540. 
Siemens-Martin process for ‘steel, 
548; regenerative furnaces for, 
549, 550. 
Silicates, artificial, 352; list of com- 
mon, 348, 355; natural, 355. 
Silicic acid from silicon fluoride, 350; 
from silicon chloride, 351; insol- 
uble after drying, 354. 
Silicic acids, composition, 354; im- 
portance of colloidal in soils, 354; 
preparation, colloidal, dialysis of, 


elec- 
436; of 


353; structure of, 354; various, 
35D: 

Silicon carbide, 349. 

Silicon dioxide, forms, 348, 351; 


formula, 354; properties, 352. 

Silicon, exercises, 368; fluoride, 
preparation, hydrolysis, 350; hexa- 
iodide, preparation, silicooxalic acid 
from, 351; occurrence, 348; prepa- 
ration, properties, 349; tetrachlo- 
ride, preparation, properties, 
hydrolysis, 351;  tetrafluoride, 
formed in etching glass, 151; 
tetraiodide, silicon hexaiodide from, 
Bo12 

Silicooxalic acid, preparation, proper- 
ties, structure, 351. 

Silicozirconates, 363. 

Silver, annual production and value, 
properties, alloys, plating, 442; 
argenticyanide, evidence of exist- 
ence of complex ions in, transfer- 
rence of ions of in electrolysis, 379 ; 


bromide, properties, conduct 
toward light, 444; use in photog- 
raphy, 445. 


Silver chloride, properties, conduct 
toward light, 444; use in pho- 
tography, 445; reduction by mer- 


cury, 441; solubility product of, 
aren 
Silver chloroplatinate, 566; dep- 


osition from potassium argenti- 
cyanide, 321; exercises, 450; for 
mirrors, 487. 
“Silver from Clay,” 
aluminium, 495. 
Silver hydroxide, 442; ionization of, 
443; hydroxide, temporary forma- 
tion, 393; iodide, properties, con- 


applied to 


596 


duct toward light, 444; iodide, use 
in photography, 445; molten, ab- 
sorption of oxygen by, 443. 

Silver nitrate, preparation, proper- 
ties, uses, 444; theory of reaction 
with salt, 376; use in determining 
the atomic weight of chlorine, 131. 

Silver nitrite, preparation, use, 444; 
oxide, dissociation pressure, 443; 
formation, 442 ; orthoarsenite, 259 ; 
peroxide, formation, 443; pyro- 
phosphate, 253. 

Silver, occurrence, metallurgy, Patti- 
son’s process, 439; cupellation, 
Parke’s process, 440; amalgama- 
tion process, electrolytic process, 
cyanide process, 441. 

Silver recovered from copper, 430, 
439; separation from lead, 439, 
440; sulfide, conversion to chloride 
in ores, 441; sulfide on coin, test 
for sulfur, 409; sulfate, prepara- 
tion, use, 444; trinitride, failure 
to react with iodine trinitride, 224. 

Sizes for paper, 338. 

Slag, blast furnace, use for cement, 
454, 543. 

Slimes, from electrolytic refining of 
copper, 430. 

Smalt, 558. 

Smaltite, 557. 

Smith, Alex., formula of calomel, 489. 

Smith, E. F., potassium fluotantalate, 
524; preparation of metallic cal- 
cium, 452. 

Smithells, separated Bunsen flame, 
301. 

Smithsonite, 481. 

Smokeless powder, 419; gun cotton 
in, 338. 

Soaking pits, in steel manufacture, 548 

Soap, manufacture, use, 332; water 
glass in, 353; soft, from lye of 
wood ashes, 414. 

Soapstone, metasilicate, 355. 

Soda, baking, 412; washing, 411. 

Soda lime, use in preparing methane, 
286; to absorb.carbon dioxide and 
water, 6. 

Sodamide, base in liquid ammonia, 
208; hydronitric acid from with 
nitrous oxide, 223; preparation, 
properties, ionization, use in mak- 
ing indigo, 410. 

Soda water, 309. 

Soddy, discovery of helium from 
radium, 473. 

Sodium aluminate, hydrolysis, 499; 
decomposition by carbonic acid, 
496; aluminate, manufacture from 
clay, 496; amalgams, composi- 





INDEX 


tion, formula for, 487; ammonium 
phosphate (microcosmic salt), 253; 
antimonite, 265; atoms, rate of 
vibration, 426. 

Sodium bicarbonate, formation from 
carbonic acid, 360; loss of carbon 
dioxide on boiling solution, 376; 
Sein or ammonia soda process, 

22 

Sodium bismuthate, oxidation of 
manganese by, 269; bisulfate, 408; 
bisulfite, 408;  bromate, 144; 
bromate, use with potassium iodide 
to illustrate strength of acids, 386. 

Sodium carbonate, by-product in 
making aluminium oxide from 
clay, 496; decomposition by acids, 
theory, 375; formation from car- 
bonic acid, 310; hydrolysis of, 385; 
Leblane soda process, 411; mono- 
hydrate, 411; transition point, 
dekahydrate, 412; uses, 412. 

Sodium chloride, crystallization, con- 
centration of solution of, 405; elec- 
trolysis of, 401; localities for, 398 ; 
properties, essential in diet, 406; 
sodium bicarbonate from, 412; 
sources, 404; theory of reaction 
with silver nitrate, 376. 


Sodium cobaltinitrite, 558; copper 
orthophosphate, 253; formate, 
oxalic acid from, 329; hydrosul- 


fide, formation, properties, 409. 

Sodium hydroxide, Castner-Kellner 
process for, 402; from sodium 
carbonate, 401; from sodium 
chloride, 401; properties, density 
of solutions, table, 403. 

Sodium hypobromite, 143 ; hypophos- 
phite, formed in preparing phos- 
phine, 248; hyposulfite, prepara- 
tion, uses, 408; hyposulfite, prep- 
aration, 186; hyposulfite, use for 
reduction of indigo, 341. 

Sodium ‘‘ hyposulfite,’’ old name 
for thiosulfate, 408. 

Sodium hypophosphate, acid, 254; 
iodide, separation of ions by cen- 
trifugal force, 114; manganate, 
preparation conversion to perman- 
ganate, 536; manufacture for 
aluminium, 495; manufacture of 
sodium peroxide from, 404; meta- 
phosphate, formation, use, 253; 
metaphosphate, from microcosmic 
salt, use, 423. 

Sodium nitrate, acid sodium sulfate 
from decomposition of, 408; for- 
mation in soil, 199; reduction to 
nitrite, 218; source, properties, 
use, 410. 


INDEX 


Sodium nitrite, preparation, 218; 
preparation, use, 410. 

Sodium, occurrence, 398; metallurgy, 
properties, 399; uses, 400; oxide, 
preparation, properties, 404; per- 
borate, bleaching by, 367; perchlo- 
rate, use in preparing perchloric 
acid, 128. 

Sodium peroxide, hydrogen peroxide 
from, 84; preparation of oxygen 
from, 21; preparation, properties, 
hydrate, uses, 404. 

Sodium phosphites, 248; pyro- 

borate, 367; pyrosulfate, prepara- 
tion, 186; pyrosulfate, prepara- 
. tion, use in analysis, 408; ses- 
quicarbonate, occurrence, 398; 
silicates, manufacture, use, 413; 
silicate, preparation, properties, 
uses, 353; sulfarsenate, 261. 

Sodium sulfate, acid, preparation, 
uses, pyrosulfate from, 408; manu- 
facture, hydrate, transition point for 
hydrate, 406 ; solubility curve, 407 ; 
constituent of Hunyadi water, 408. 

Sodium sulfide, preparation from 
sodium hydroxide, from sodium 
carbonate and sulfur, as test for 
sulfur, 409; sulfite, acid, use in 
cider, 175; sulfite, acid, prepara- 
tion, uses, 408. 

Sodium stannate, use in fireproofing 
cotton goods, 513; stannite, prep- 


aration, reducing agent, 510; 
tetrathionate, from sodium thio- 
sulfate, 187; tetrathionate, for- 


mation in iodine titrations, 409. 
Sodium thiosulfate, conversion to 


tetrathionate, 187; preparation, 
use in photography, 408; in 
lixiviation processes, 409; anhy- 


drous, 409; precipitant for copper, 
409; structure, 409; use in 
photography, 445. 

Sodium trinitride, formation, 410; 
trinitride, preparation, 223; tung- 
state, 531; zincate, 483. 

Softening water, by boiling, 463; by 
use of sodium salts, 464; Clark’s 
process for, 464; ‘‘ permanently ” 
hard waters with sodium carbonate, 
phosphate, fluoride or borate, 311. 

Soils, formation of, 494; formation 
of nitrates in, 199; importance of 


colloidal silicic acids in, 353; po- 
tassium in, 414. 

Solder, 509. 

Soluble glass, 353; manufacture, 
uses, 413. 


Solubility of salts, general statement, 
. 393; graphical representation, 80. 





597 


Solubility product, 377; rule for, 378; 
relation to solubility of magnesium 
hydroxide, 479; relations for 
tan and bi-bivalent salts, 

78. 

Solute, defined, 79. 

Solution, defined, 79; formular, def- 
inition of, 183; molar, definition 
of, 183; normal, definition of, 184; 
standard, definition of, 183. 

Solution pressure, defined, 435. 

Solutions, chemical activity in, 81; 
differences in potential between, 
437; in liquid ammonia, 207; 
supersaturated, 80. 

Solvay, discovery of ammonia soda 
process, 400. 

Solvent, defined, 79. 

Sommerfield, quantum theory, 398. 

Specific heat, fixes atomic weight of 
indium, 506; of elements, 397; 
of gases at constant volume and 
constant pressure, 236. 

Spectra, comparison of, 427; 
gases, Pliicker tubes, 427; 
metals, how obtained, 427. 

Spectroscope, direct vision, 427; de- 
scription of, 425. 

Spectrum analysis, 424. 

Spectrum, continuous, bright and 
dark lines, 425; theory of dark 
line, diffraction, solar, 426. 

Speed of chemical reactions, 149; 
relation to chemical affinity, 149; 
and concentration, 149. 

Speed relation of two reactions 
calculated at equilibrium, 151. 

Sphalerite, 161, 481. 

Stafford, molecular weight of sulfur 
vapor, 163. 

Standard solutions, definition of, 184. 

Stannate, use in fireproofing cotton 
goods, 513. 

Stannic acids, table of, 511. 

Stannic acid, preparation, properties, 
511. 

Stannic chloride, preparation, con- 
duct in solution, 512; oxide, prep- 
aration, 510; use in glass, 511; 
sulfide, 512. 

Stannous chloride, preparation, uses, 
reducing agent, 510; oxide, 510; 
sulfide, 510; sulfide dissolved by 
ammonium polysulfide, 422. 

Stannum, 11. 

Starch, glucose from, 334; source, 
manufacture, varieties, 335; ap- 
pearance of granules, cooking of, 
336. 

Stassfurt, potassium chloride from, 
415 


of 
of 


598 INDEX 


Stearic acid, source, 331. 171; of Group II, solubility, 491; 
Stearin, 331. roasting in metallurgy, 391; theory 
Steel, cementation, analyses, 551; of precipitation by alkaline sul- 


alloy, 552; cast, tempering, 545; fides, 170. 

theory of tempering, 546; Besse-| Sulfites, as germicides, 174; prep- 
mer, 547; acid and basic Bessemer, aration and uses, 175. 

open hearth, Siemens-Martin, 548 ; | Sulfocarbonates, formed from carbon 
resistance to corrosion increased by bisulfide, 317. 


copper, 431. Sulfocarbonic acid, formation, de- 
Stellar photometry, use of selenium composition, 318. 

in, 190. Sulfocyanates, see Thiocyanates. 
Stereotype metal, 264, 269, 515. Sulfur, allotropic forms of, 162; 
Stewart, standard of ventilation, 231. amorphous, 162; boiling point, 
Stibine, 264. 1633. burning in 9 Oxygen, e2a- 
Stibium, 11. by Chance process, 160; com- 
Stibnite, 263. pounds containing halogens, 188.. 
Stockholm, holmium named for, 505.| Sulfur dioxide, as germicide and 
Stokes, ferric sulfide, 556. disinfectant, 173; bleaching by, 
Storage batteries, theory of, 516. 173% catalysis of conversion to 
Storage of gases, 22. sulfuric acid by oxides of nitrogen, 
Strass, 467. 178; coefficient of expansion, 38; 
Strength of acids, definition, 168; deviation from’ Boyle’s law, : 35; 

illustration, 386. from burning sulfur, 23; from iron 
Strength of organic bases, 339. pyrites, 177; preparation by burn- 
Strong acids and bases defined, 386. ing sulfur, 172; preparation by 
Strontianite, 468. reduction of sulfuric acid, 173; 
Strontium carbonate, 468; com- preparation from acid sodium 


pounds, 468; flame color, 471; sulfite, 173; properties, 173; solu- 
hydroxide, 468; nitrate, use in tion in water, 174; uses, 173. 
fireworks, 468; occurrence, 467;}| Sulfur, exercises, 196; family, table 
sulfate, 468. of compounds, 192; flowers of, 
Structural formulas, basis for, 323. 160; gaseous, Ss, S. and S, 163% 
Structure of compounds, relation to group, 160; heat of combustion, 27. 
valence, 65; of the oxyacids of | Sulfur hexafluoride, 188. 


chlorine, 130. Sulfur, in Louisiana, getting Of 161; 
Strychnine, 343; salt of chloroauric in petroleum removed with copper 

acid, 450. oxide, 289; in proteins, 343; in 
Strychnos nux vomica, strychnine Sicily, getting of, 160. 

from, 343. Sulfur, lime-, wash, 164; mobile 
Study of chemistry, 18. liquid (Sq), 162; monochloride, 
Subdivisions of science, 3. preparation, uses, 188; monoclinic, 
Substance, definition, 7. 162; native, source of, 160; 
Substantive dyes, 342. occurrence, 160. 
Substitution, 287 ; use in determining | Sulfur, production in U. S. and in- 

structure, 324. world, 161; properties, 163 ; rhom- 
Sudbury, nickel from, 559. bic, 162; roll brimstone, 161; test 
Sugar, cane, occurrence, manufac- for with sodium carbonate on 


ture, properties, effect on polarized charcoal, 409. 
light, 333; hydrolysis, 334; puri-| Sulfur trioxide, absorption by con- 
fication by bone black, 278; solu- centrated sulfuric acid, 176; for- 


tions clarified by basic lead acetate, mation from sulfur and oxygen 
519; use of strontium hydroxide reversible reaction, 175;  poly- 
in manufacture of, 468. meric, 176; preparation, 175; 
Sugar of lead, 519. properties, 176; sulfuric acid 
Sulfantimonates, 268. from, 45. 
Sulfantimonites, 268. Sulfur, uses, 164; viscous liquid (Sy), 
Sulfarsénates, formation, 261. 163. 
Sulfarsenites, formation, 261. Sulfuric acid as dehydrating agent, 
Sulfates, list of insoluble, 183. 182; ‘‘ chamber acid,”’ 180; 
Sulfides, basis of groups of qualita- chamber process for, 177: con- 


tive analysis, 166; hydrolysis of, centration of, 180; directions for 


INDEX 


dilution of, 182; dissociation of, 
180; electrolysis of, 9, 47; ‘‘ fum- 
ing,’’ 176; fuming from ferric sul- 
fate, 556; Gay-Lussac tower, 179; 
hydrates of, 181; ionization of, 181; 
moisture left in gas by, 54; prep- 
aration by ‘‘ chamber _ process,”’ 
177 ; properties, 180; reaction with 
copper, 173 ; reduction by hydroidic 
acid, 145; reduction by hydro- 
bromic acid, 142. 

‘“* Sulfuring ”’ fruit, 164, 174. 

Sulfurous acid as reducing agent, 
174; formation, 174; ionization of, 
174; “strength’”’ of, 174; struc- 
ture, 174. 

Sulfuryl chloride, hydrolysis, 188; 
preparation, properties, 188. 

Sun’s corona, heliumin, 237; photo- 
sphere, spectrum, 426. 

Superphosphate, calcium, manufac- 
ture, use in fertilizers, 461. 

Supersaturated solutions, 80. 

Sylvite, 414. 

Symbols of elements, 11. 

Sympathetic ink, 557. 

Synthesis, definition, 66. 

Synthesis, quantitative, of water by 
volume, 66. 

Systems of crystallography, 193. 


Table of atomic weights, 10; abso- 
lute potentials of elements, elec- 
tromotive series, 436; acidity or 
alkalinity of indicators, 388; 
compounds of elements of Group V, 
271.; compounds of sulfur family, 
192; correction of apparent weight 
of water to volume, 73; correction 
of readings of barometer for glass 
and brass scale, 36; correction of 
readings of barometer for latitude 
and altitude, 37; degree of 
ionization, acids, 383; bases, 383; 
salts, 384; density and volume of 
water, 73; density of gases, 95; 
density of solutions of sodium hy- 
droxide, 403; deviation of gases 
from Boyle’s law, 35; deviation of 
gases from law of Charles, 38; dis- 
sociation of water, 61; elements in 
earth’s crust, 11; elements of sul- 
fur group, 160; equilibrium be- 
tween nitrogen and oxygen, 216; 
equilibrium of hydrogen, iodine and 
hydriodic acid, 147; groups and fam- 
ilies of elements, 371; heat of com- 
bustion of C,S, P, Fe and Hg, 27; 
halogen acids, 139; halogen family, 
139; nonmetallic elements, 348; 
Periodic System, 134, 135; proper- 


599 


ties and compounds of elements of 
Group VIII, 562; rate of decom- 
position of hydriodic acid, 148; 
rate of formation and decomposi- 
tion of hydriodic acid, 150; stannic 
acids, 511; values of calorie, 10° 
30°, 33; vapor pressure of ice and 
water, 75; vapor pressure of sys- 
tems containing calcium sulfate, 
459; varieties of ferrite, 546. 
ae preparation of hydroxylamine, 


Talc, metasilicate, 355. 

Tannic acid, mordant, 342. 

Tantalum, discovery, properties, 523, 
ey compounds, 524, electric light, 

Tartar emetic, 266. 

ra acid, structure, source, salts, 

Tasmania, tin from, 508. 

Teeth, amalgam for, 487. 

Telluric acid, 190. 

Tellurium, anomalous position in 
periodic tabie, 138; atomic weight 
of, 190; dioxide, 190; occurrence, 
preparation, compounds, 190. 

Tellurous acid, 190. 

Temperature, absolute, 39, critical, 
233; effect of on a gas, 38; inter- 
national scale of, 32; kindling, 24; 
of flames, 302; of interior of earth, 
relation of radioactivity to, 476; 
units of, 32. 

Ieee) thermometers for high, 

6. 

Temporary hardness, 310. 

Terbium compounds, 505. 

Terne plate, 509. 

Tertiary salts, 249. 

Tetragonal pyramid (crystal), 194. 

Tetragonal system (crystallography), 
194. 

Tetrahedron (crystal), 193. 

Tetrahexahedron (crystal), 193. 

Tetramethyl ammonium hydroxide, 
analogy with ammonium hydrox- 
ide, 206. 

Tetraphosphorus heptasulfide, 254. 
Tetraphosphorus trisulfide, proper- 
ties, 254; use for matches, 243. 

Tetrathionic acid, 188. 

Thallium, discovery, compounds, 507. 

Theorem of Le Chatelier, 111. 

Theory, atomic, 14; electrochemical, 
influence on formulas of minerals, 
356; the electron, 181; electron, 
in relation to metals and non- 
metals, 370; electron, relation to 
ionization, 206, 379; colloidal solu- 
tions, 262; decomposition of car- 


600 


bonates by acids, 375; chamber 
process for manufacture of sulfuric 
acid, 178; hydrolysis of salts, 385; 
neutralization, 384; the quantum, 
398; storage batteries, 576; 

Thermite process, Goldschmidt’s, 
497; chromium by, 524; tungsten 
by, 530. 

Thermodynamic scale, melting points 
Onvove: 

Thermometers, glass for, 467; of 
fused quartz, 352; mercury and 
international scales for, 486; zero 
point correction, special for high 
temperatures, 486; point on, 
fixed by transition point of sodium 
sulfate, 406. 

Thio-, prefix, 187. 

Thiocyanate, potassium, 321. 

Thiosulfates, preparation,-use, 186. 

Thiosulfuric acid, formation and de- 
composition, 187; formation from 
sulfur monochloride, 188. 

Thomsen, heat of combustion of 
hydrogen, 65. 

Thompson, J. J., the electron theory, 
181; atoms of metallic elements, 94. 

Thompson, Sir William, size of mole- 
cules, 16. 

Thorianite, 364. 

Thorium, in monazite sand, proper- 
ties, dioxide, sulfate, nitrate, use 
in Welsbach mantles, 364; series 
of elements, 475. 

Thulium, compounds, 506. 

Thyroid gland, iodine in, 144. 

Tilkerode, thallium from, 507. 

Time, units of, 32. 

Tincture, defined, 144. 

Tin, occurrence, sources, metallurgy, 
508; properties, alloys, uses, tin 
plate, 509; compounds, 510; re- 
covery from tin scrap, 508. 

Titanium, compounds of as mordants, 
363; occurrence, properties, com- 
pounds, 362; oxide, solution in 
sodium pyrosulfate, 408; separa- 
tion from silica, detection, 362; 
tetrafluoride, 362; test for hydro- 
gen peroxide, 86. 

Tolman, separation of ions by cen- 
trifugal force, 114. 

Toluene, 283. 

Toning in photography, 445. - 

Tool steel, high-speed, 530. 

Topaz, 349. 

Torpedoes, gun cotton in, 338. 

Tourmaline, 349. 

Toxins, 344. 

Transition or quadruple point, for 
sodium sulfate, 406. 


INDEX 


Transition points, for steel and ferrite, 
546. 

Trautz, theory of sulfuric acid manu- 
facture, 179. 
Triammonium dodekamolybdate, 529. 

Tribasic acids, defined, 183. 

Tricalcium phosphate, 249. 

Triclinic system (crystallography), 
196. 

Tridymite, 352. 

Trimethyl amine, 204. 

Triphosphorus hexasulfide, 254. 

Triple-effect evaporation, 405. 

Triple point, definition, 78. 

Triple point water-ice-water-vapor 
above 0°, 407. 

Trisilver phosphate, 253. 

Trisodium phosphate, 249; aliza- 
rine green as indicator for, 251. 

Trisilicates, 356. 

Trisilicic acids, 355. 

Trisodium phosphate, hydrolysis, al- 
kaline reaction of, 252. 

Trithionic acid, 188. 

Trivalent, definition, 64. 

Trypsin, 344. 

Tuberculosis, relation to ventilation, 
231). 

Tungsten, history, preparation, prop- 
erties, use in lamps, in tool-steel, 
530. 

Tungsten bronze colors, 531. 

Turmeric paper, test for boricacid, 368. 

Turnbull’s blue, 321. 

Tuyéres of blast furnace, 541. 

Type metal, 264, 175. : 

Typhoid fever from impure water 
supply, 83. 


Ultramarine, natural and artificial, 
502. 

Ultra-violet light, use in purifying 
water, 83. 

Uni-bivalent salts, law of solubility 
product not general for, 378. 

Unimolecular reactions, 150. 

Unit for atomic weights, 68; electri- 
cal charge, 4388; of length, 31; 
of power, 33; of temperature, 32; 
of volume, 31; of weight, 31. 

Units, absolute, 33; electrical, 33; 
of energy, 32; of heat, 33; of 
mechanical energy, 33; of time, 32. 

Uni-univalent salts, law of solubil- 
ity product for, 378. 

Univalent, definition, 64. 

Univariant, definition, 78. 

Unsaturated compounds, definition, 
291. 

Uraninite, 531; helium in, 237. 

Uranium chlorides, 532; catalyzer 


INDEX 


for synthesis of ammonia, 201; 
occurrence, 531, properties, com- 
pounds, 532; radium from, 473, 
half-life, 474; series of radioac- 
tive elements, 475; sulfate, 532. 

Uranyl acetate, 532; compounds, 
532; nitrate, 532. 

Urea, formed in body and from am- 
monium cyanate, 345. 

Use of Indicators, 387 ; for weak acids 
and bases, 389. 


Vacuum desiccator, 84; high, by 
means of charcoal, 278. 


Valence, definition, 63;  illustra- 
tion, 63; of elements in oxides 
and salts, 157; of elements in 


periodic system, 133; relation to 
equivalents and Faraday’s law, 
438; use in determining structure, 
323; use in writing equations, 156. 

Vanadinite, 522. 

Vanadium, occurrence, 
uses, compounds, 522. 

Vanadous compounds, 522. 

van’t Hoff, definition of osmotic 
pressure, 360. 

van’t Hoff-Le Chatelier, principle of, 
111; applied to reversible reac- 
tion, 148; applied in the prepara- 
tion of sulphur trioxide, 175; 
epplicd in synthesis of ammonia, 


properties, 


Vapor pressure, definition, 74; of 
hydrates, 82; of ice and water, 75. 

Vaselin, 290. 

Bacquelin, discovery of chromium, 

Vegetable foods, 347. 

Venetian red, 555. 

Ventilation, 230 ; standardof, 231; lack 
of causes disease, 231. 

Vermilion, 489. 

Vinegar, 329. 

Vitriol, blue, 433 ; oil of, 46, 434; defi- 
nition, 434; green, 554; white, 483. 

Volatility, effect on reactions, 374. 

Be pances, source of carbon dioxide, 

29. 
Volt, defined, 33. 
Volume, unit of, 31. 


Washing soda, 411. 

Water, as a solvent, 79; calculation 

' of the composition by weight, 68; 
of crystallization, 82; degree of 
ionization, 383; density of at 
different temperatures, 73. 

Water, determination of composition 
by copper oxide, 69; by weighing 
oxygen and hydrogen, 71. 


601 


Water, determination of, in air by 
weighing and dew point, 232. 

Water, dissociation of, 59; effect of 
on chlorides, 112; filters, charcoal 
not efficient in, 278; heat of fusion, 
74; heat of vaporization, 74. 

Water gas, heat relations in manu- 
facture, 298; percentage composi- 
tion, 299; enriched, 297;  per- 
centage composition, 299, 296; 
carbon monoxide in, 297. 

Water glass, 353; of hydration, 82; 
hydrolysis of chlorides by, 115; 
ionization of, 171, 383; maximum 
density, effect, 72; phases and 
triple point, 78; properties of, 72; 
purification of, 83. 

Water, qualitative analysis and syn- 
thesis of, 66; quantitative ioniza- 
tion of in relation to indicators, 387 ; 
sea, amount of carbon dioxide in, 
230. 

Water, use of aluminium sulfate in 
purifying, 500. 

Water vapor, effect of, on the volume 
of a gas, 76. 

Water, vapor pressure of, 75; table, 
75; weight of 1 liter at different 
temperatures, 73. 

Water-soluble phosphoric acid, 461. 

Waters, effervescent, 309; hard, 310; 
natural, 82; radioactive, mineral, 
476. 

Watt defined, 34. 

Waves, explosion, 301. 

Weak acids and bases, defined, 386. 
Weber, preparation of chloroplatinic 
acid, 566. f 
Weight and mass, relation, 32; unit 

Ofnals 

Weights and Measures, International 
Bureau of, 31. 

Welding by thermite process, 498; 
use of borax in, 367. 

Weldon process, 103. it 

Welsbach, resolution of dydimium, 
504. 

Welsbach mantles, manufacture, 364 ; 
theory of, 365. 2 ’ 
Wentzki, theory of sulfuric acid 
manufacture, 179. ; : 
Werner, formula for ammoniocupric 
sulfate, 434; periodic table, 138; 
theory of isomeric hydrates of 
chromic chloride, 526; theory of 

valence, 559. 

Weston cell, electromotive force, 437, 
438. 

Whisky, 325. ! 

White lead, manufacture, properties, 
520; poisoning by, 521. 


602 


White vitriol, 483. 

Willemite, 481. 

Wilson, discovery of scandium, 503. 

Winkler, discovery of germanium, 
361% 

Witherite, 468. 

Wohler, discovery of aluminium, 495; 
synthetic urea, 345. 

Wolfram, 11. 

Wolframite, 530. 

Wood alcohol, 325; charcoal, 277; 
oak, composition, 280. 

Wood spirit, 324. 

Wood’s metal, 269. 

Woolrich, Sherardized iron, 482. 

Wrought iron, development of 
puddling process, 544;  replace- 
ment by other irons, 545. 


Xenon, discovery, 238. 
X-rays, 474. 
Xylene, 283. 


Yaryan evaporator, 405. 

Yeast, fermentation of sugar by, 325. 

Ytterbium, compounds, 504. 

Yttrium group of rare earths, 503; 
occurrence, compounds, 503. 


Zechentmayer, ferrous chloride and 
nitric oxide, 554. 
Zero, absolute, 40. 


INDEX 


Zero group, 236. 

Zero point, corrections for ther- 
mometers, 486; 
thermometers, 467. 

Zinc chloride, preparation, properties, 
use as wood preservative, 483; 
compounds, effect of ammonium 
hydroxide on solutions of, 491; 
Pe asthe formation, amphoteric, 

Zinc, in brass and bronze, 431; oxide, 
green with cobalt nitrate, 558. 

Zinc, occurrence, metallurgy, proper- 
ties, uses, galvanized iron, 481; 
preparation, properties, uses, 482; 
Sherardized iron, 482. 

Gee 483; use in Clark cell, 


Zinc sulfide, fluorescence of, with ' 


radioactive elements, 472; forma- 
tion, conduct toward acids, 483; 
solubility, 491. 

Zinc, use in Parke’s process, alloys 
with lead, 441. 

Zincic acid, salts of, 483. 

Zirconates, 363. 

Zirconium dioxide, use in Nernst 
lamp, oxyhydrogen light and for 
crucibles, 363. 

Zirconium, occurrence, 
compounds, uses, 363. 

Zymase, 344. 


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depression of in 


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